Chemistry Check Off List:



NAME:______________________________________________ Period ____

Core Curriculum Check-Off List

Regents Chemistry

The following information is everything you need to know and be able to do to attain Mastery of the Regents Chemistry Curriculum. The bold, underlined words are important vocabulary words that you should be able to define and use properly in explanations. This is a study guide for what you will be tested on throughout the year. The objectives are divided into categories of “Knowledge” (what you have to know) and “Application” (what you have to be able to do). Check off each objective when you fully understand it. If you do not understand an objective, ask questions before you are tested on it.

|Mathematics Analysis & Graphing |

|1. |Identify independent and dependent variables in an experiment and correctly plot them on an axis |

| | |

| |Example hypothesis: Chemistry students who do their homework will have higher test scores than students who do not do their homework. |

| |( X-axis (horizontal): the independent variable is the one that is manipulated by the experimenter. (“The one I change.” – do homework/not do |

| |homework) |

| |( Y-axis (vertical): the dependent variable is the one that changes based on the independent variable. (The data you collect – test scores) |

|2. |Express uncertainty in measurement by properly using significant figures |

| |Identify the number of sig. figs. in data |

| |Round to the correct number of sig. figs. on calculations: |

| |( Addition and Subtraction (round to least precise [farthest left] place value) |

| |( Multiplication and Division (round to the lowest # of digits in data) |

| |( Combined Rules (Add/Subtract first, then Multiply/Divide) |

|3. |Identify relationships between variables from data tables and graphs (direct or inverse relationships) |

|4. |Understand what is meant by conditions of Standard Temperature and Pressure (STP) |

| |(Table A) |

|5. |Recognize and convert between units on various scales of measurement (Tables C, D, and T) |

| | |

| |( Temperature: Celsius ↔ Kelvin ( Length: meters ↔ centimeters ↔ millimeters |

| |( Mass: grams ↔ kilograms ( Pressure: kilopascals ↔ atmospheres |

| |( Thermal Energy: joules ↔ kilojoules ( Amount of Substance: |

| |GFM = 1 mole = 6.02x1023 particles |

| |grams ↔ moles ↔ atoms or molecules |

|6. |Use the density equation on Table T to solve for density, mass, or volume, given the other two values |

|7. |Calculate percent error (Table T) |

|Atomic Concepts |

| |Knowledge |Application |

|1. |Atoms are the basic unit (building block) of matter. |Explain what happened during the gold-foil experiment and what it |

| |Atoms of the same kind are called elements. |showed. |

| |The modern model of the atom has developed over a long period of time |Gold foil was bombarded (hit) with positively charged alpha |

| |through the work of many scientists. |particles. Most alpha particles passed through the gold foil, but |

| | |some were deflected. This showed that: |

| | |1. the atom is mostly empty space |

| | |2. the nucleus is small, positively charged, and |

| | |located in the center of the atom |

|2. |The three subatomic particles that make up an atom are protons, neutrons, and electrons. |

| |The proton is positively charged, the neutron has no charge, and the electron is negatively charged. |

| |(This is also referenced on Table O. Check it out!) |

|3. |Each atom has a nucleus with an overall positive charge, made up of protons |Determine the nuclear charge of an atom. (equal to the number of |

| |and neutrons. |protons in the nucleus) |

| |The nucleus is surrounded by negatively charged electrons. | |

|4. |Atoms are electrically neutral, which means that they have no charge (# |Determine the number of protons or electrons in an atom or ion when |

| |protons = # electrons) |given one of these values. (Periodic Table) |

| |Ions are atoms that have either lost or gained electrons and are either |Compare the atomic radius and ionic radius of any given element |

| |positively or negatively charged |Ex: A chloride ion has a larger radius than a chlorine atom because |

| |When an atom gains one or more electrons, it becomes a negative ion and its |the ion has an extra electron. A sodium ion has a smaller radius |

| |radius increases. |than a sodium atom because the ion has lost an electron. |

| |When an atom loses one or more electrons, it becomes a positive ion and its | |

| |radius decreases. | |

|5. |The mass of each proton and each neutron is approximately equal to one |Calculate the mass of an atom given the number of protons and |

| |atomic mass unit (AMU). |neutrons |

| |An electron is much less massive (has almost no mass) compared to a proton |Calculate the number of neutrons or protons, given the other value |

| |or neutron. | |

|6. |In the modern model of the atom, the WAVE-MECHANICAL MODEL (electron cloud), the electrons are in orbitals (clouds), which are defined as regions|

| |of most probable electron location. |

|7. |Each electron in an atom has a specific amount of energy. |

| |Electrons closest to the nucleus have the lowest energy. As an electron moves away from the nucleus, it has higher energy. |

|8. |Electron configurations show how many electrons are in each orbital. |Distinguish between ground state and excited state electron |

| |When all of an atom’s electrons are in the orbitals closest to the nucleus, |configurations. (Be careful to keep the same number of electrons |

| |the electrons are in their lowest possible energy states. This is called |when writing the excited state.) |

| |the ground state. | |

| |When an electron in an atom gains a specific amount of energy, the electron | |

| |is at a higher energy state (excited state) | |

|II. Atomic Concepts (continued) |

| |Knowledge |Application |

|9. |When an electron returns from a higher energy (excited) state to a lower | |

| |energy (ground) state, a specific amount of energy is emitted. This emitted|Identify an element by comparing its bright-line spectrum to given |

| |energy can be used to identify an element. |spectra |

| |The flame test is an example of the bright-line spectrum visible to the | |

| |naked eye. The color can determine the identity of a positive ion in a | |

| |compound. | |

|10. |The outermost electrons in an atom are called the valence electrons. In |Draw a Lewis electron-dot structure of an atom. |

| |general, the number of valence electrons affects the chemical properties of |Distinguish between valence and non-valence electrons, given an |

| |an element. |electron configuration |

|11. |Atoms of an element that contain the same number of protons but a different |Calculate the number of neutrons in an isotope of an element given |

| |number of neutrons are called isotopes of that element. |the isotope’s mass |

|12. |The average atomic mass of an element is the weighted average of the masses |Given an atomic mass, determine the most abundant |

| |of its naturally occurring isotopes. |isotope |

| | |Calculate the atomic mass of an element, given the masses and |

| | |abundance of naturally occurring isotopes |

|The Periodic Table |

| |Knowledge |Application |

|1. |The placement or location of an element on the Periodic Table gives an indication of |Explain the placement of an unknown element in the |

| |physical and chemical properties of that element. |periodic table based on its properties |

| |The elements on the Periodic Table are arranged in order of increasing atomic number. | |

|2. |The number of protons in an atom (atomic number) identifies the element. This goes at the|Interpret and write isotopic notations. |

| |bottom left corner of the symbol for that element. |Ex: |

| |The sum of the protons and neutrons in an atom (mass number) identifies an isotope. The |C-12, C-13, and C-14 are isotopes of the element |

| |mass number is placed at the top left corner of the symbol for an element OR is placed |carbon |

| |after the element symbol or name and a dash. | |

| |Ex: Three different ways to write carbon with a mass number of 14: | |

| |C, carbon-14, or C-14 | |

|3. | |Identify the properties of metals, metalloids, |

| |Elements are classified by their properties, and located on the periodic table as metals, |nonmetals and noble gases |

| |metalloids, nonmetals, or noble gases. |Classify elements as metals, metalloids, nonmetals, or|

| | |noble gases by their properties |

|III. The Periodic Table (continued) |

| |Knowledge |Application |

|4. |An element’s atomic radius, first ionization energy, and electronegativity determine its physical and chemical properties |

|5. |Substances can be differentiated by their physical properties. |Identify and give examples of physical properties |

| |Physical properties of substances include melting point, boiling point, density, |Describe the states of the elements at STP (solid, |

| |conductivity, malleability, solubility, and hardness. |liquid, or gas). (Table S) |

|6. |Substances can be differentiated by chemical properties. |Identify and give examples of chemical properties |

| |Chemical properties describe how an element behaves during a chemical reaction and include|Describe the difference between physical and chemical |

| |reactivity, flammability, and toxicity. |properties of substances |

|7. |Some elements exist as two or more forms in the same phase. These forms differ in their molecular or crystal structure and therefore in their |

| |properties. The word to describe this phenomenon is ALLOTROPE. |

| |Ozone and oxygen gases are allotropes of each other. Ozone is O3 and it is very dangerous to our health. Oxygen gas is O2 and we need it to |

| |survive. |

| |Diamonds and graphite (better known as pencil lead) are both forms of the element carbon. They have different molecular structures and very |

| |different properties. |

|8. |For Groups (also called families) 1, 2, and 13-18 on the Periodic Table, elements within |Determine the group of an element, given the chemical |

| |the same group have the same number of valence electrons (helium is the exception) and |formula of a compound |

| |therefore similar chemical properties. |Ex: A compound has the formula XCl2, element X is in |

| |Elements in the same Period (row) have the same number of principal energy levels (shells)|Group 2 |

| |which contain electrons. |Determine the number of energy levels containing |

| | |electrons given an element’s Period and vice versa |

|9. |The succession of elements within the same GROUP (top to bottom) demonstrates |Compare and contrast properties of elements within a |

| |characteristic trends: differences in atomic radius, ionic radius, electronegativity, |group or a period for groups 1, 2, and 13-18 on the |

| |first ionization energy, and metallic/nonmetallic properties. |periodic table |

|10. |The succession of elements across the same PERIOD (left to right) demonstrates |Understand and be able to explain the trends in terms |

| |characteristic trends: differences in atomic radius, ionic radius, electronegativity, |of nuclear charge and electron shielding |

| |first ionization energy, and metallic/nonmetallic properties. | |

|Matter, Energy, & Change |

| |Knowledge |Application |

|1. |Matter is anything that has mass and volume (takes up space). |Identify specific examples of matter as an element, |

| |Matter cannot be created nor destroyed, only transformed. |compound, or mixture |

| |Matter is classified as a pure substance (element or compound) or as a mixture of | |

| |substances. | |

|IV. Matter, Energy, & Change (continued) |

| |Knowledge |Application |

|2. |Energy is not matter (does not have volume) |Distinguish between matter and energy |

| |Energy can exist in different forms, such as kinetic, potential, thermal (heat), sound, | |

| |chemical, electrical, and electromagnetic. | |

| |Energy cannot be created or destroyed, only transformed. | |

|3. |During a physical change, particles of matter are rearranged. Examples of physical |Differentiate between physical and chemical changes in |

| |changes include freezing, melting, boiling, condensing, dissolving, crystallizing, and |matter |

| |crushing into a powder |Identify and give examples of physical changes and |

| |During a chemical change, NEW substances are formed with new properties. Examples of |chemical changes in matter |

| |chemical changes include combustion (burning), rusting, and neutralizing an acid or | |

| |base. | |

| |Energy can be absorbed or released during physical and chemical changes. | |

|4. |A pure substance (element or compound) has a uniform composition and constant properties|Draw and interpret particle diagrams for elements, |

| |throughout a given sample, and from sample to sample. |compounds, and mixtures |

| |Mixtures are composed of two or more different substances that can be separated by | |

| |physical means. | |

|5. |Elements are substances that are composed of atoms that have the same atomic number. Elements cannot be broken down by chemical change. |

|6. |A compound is a substance composed of two or more different elements that are chemically|Describe differences in ionic and molecular/covalent |

| |combined in a fixed proportion. |compounds |

| |A compound can be broken down by chemical means, such as during a chemical reaction. |Identify a compound as ionic or molecular/covalent |

| |Two major categories of compounds are ionic and molecular (covalent) compounds. |compound given its properties |

| |A chemical compound can be represented by a specific chemical formula and assigned a |Name compounds based on their chemical formulas |

| |name based on the IUPAC system. |Determine the formula of a compound given its name |

|7. |When different substances (elements or compounds) are mixed together and do NOT |Interpret particle diagrams as showing homogeneous or |

| |chemically react, a mixture is formed. |heterogeneous mixtures |

| |The amounts of substances in a mixture can vary. Each substance in a mixture retains |Give examples of homogeneous and heterogeneous mixtures |

| |its original properties. | |

| |The composition of a mixture can vary. If the substances are uniformly (evenly) | |

| |distributed throughout the mixture, it is called a homogenous mixture. If the | |

| |substances are unevenly distributed, it is called a heterogeneous mixture. | |

|8. |Differences in properties such as density, particle size, molecular polarity, boiling |Describe the processes of filtration, distillation, and |

| |point, freezing point, and solubility allow physical separation of the components of the|chromatography and the types of mixtures they are used to |

| |mixture. |separate |

|Chemical Bonding |

| |Knowledge |Application |

|1. |Atoms bond with other atoms to gain a stable electron |Determine the noble gas configuration an atom will achieve when bonding |

| |configuration. | |

| |Noble gases are already stable and tend to not bond/react.| |

|2. |When a chemical reaction takes place, existing bonds must be broken in order for new bonds (and new compounds) to be formed. |

| |When a bond is broken, energy is absorbed. When a bond is formed, energy is released. |

|3. |Electron-dot diagrams (Lewis structures) are used to |Draw Lewis dot structures for any given element, ion, or compound |

| |represent the valence electron arrangement in elements, | |

| |ions, and compounds. | |

|4. |Chemical bonds are formed when valence electrons are: |Identify the type of bonding in a compound (ionic or covalent), given the elements that |

| |( transferred from one atom to |make it up |

| |another (ionic) |Demonstrate bonding concepts using Lewis Dot Structures (electron-dot diagrams) for |

| |( shared between atoms (covalent) |ionic and covalent compounds |

| |( mobile within a metal (metallic) |Ionic compounds – after the transfer of electrons, the positive ion should have no dots.|

| |Metals tend to react with nonmetals to form ionic bonds. |The negative ion should have 8 dots around it. Put brackets around the ions. Check the|

| |Nonmetals tend to react with other nonmetals to form |periodic table to find the charge associated with the ion and place this charge outside |

| |molecular (covalent) bonds. |of the brackets. *Be sure to include coefficients if there are more than one of the |

| |Ionic compounds containing polyatomic ions have both ionic|same kind of ion.* |

| |AND covalent bonds. |Covalent compounds – each atom in a covalent compound must end up with 8 dots around it |

| | |– except for hydrogen (only 2 dots). *If you use lines, remember that one line |

| | |represents 2 electrons being shared.* |

|5. |In a multiple covalent bond, more than one pair of |Draw electron-dot diagrams and give examples of molecules with multiple covalent bonds |

| |electrons is shared between two atoms. | |

|6. |Electronegativity indicates how strongly an atom of an |Distinguish between and give examples of nonpolar covalent bonds and polar covalent |

| |element attracts electrons in a bond. |bonds |

| |The electronegativity difference between two bonded |If two atoms of the same element share electrons, the bond is nonpolar (ex: H–H) |

| |determines the degree of polarity in the bond. |If atoms of two different elements share electrons, the bond is polar (ex: H–Cl) |

|7. |Molecular polarity is determined by the distribution of |Distinguish between bond polarity and molecular polarity |

| |electrons in a covalent compound. |Draw Lewis Dot Structures for all of the compounds listed to the left, including the |

| |Symmetrical distribution of electrons results in |diatomic elements |

| |(nonpolar) molecules (Ex: CO2, CH4 and all diatomic |Determine whether a molecule is polar or nonpolar, given its structure |

| |elements) | |

| |Asymmetrical distribution of electrons results in (polar) |*SNAP* |

| |molecules (Ex: HCl, NH3, H2O) | |

|Chemical Formulas, Reactions & Stoichiometry |

| |Knowledge |Application |

|1. |Chemical formulas are used to represent compounds. |Determine the empirical formula from a molecular |

| |The main types of chemical formulas include: empirical, molecular, and structural. |formula |

| |An empirical formula is the simplest whole-number ratio of atoms in a compound. |Draw structural formulas for covalent (molecular) |

| |Molecular formulas are chemical formulas that show the actual ratio of atoms in a molecule of |compounds |

| |that compound. | |

| |Structural formulas can also be used to represent covalent compounds. These use lines to show | |

| |covalent bonds between atoms and also show the geometrical arrangement of atoms in the compound.| |

|2. |One mole of any substance is equal to 6.02 x 1023 pieces of that substance. |Calculate the molar mass (gram-formula mass) of a |

| |The formula mass of a compound is equal to the sum of the atomic masses of its atoms (units are |substance |

| |atomic mass units) |Determine the molecular formula, given the |

| |The molar mass (gram-formula mass) of a substance is equal to the formula mass in grams – hence |empirical formula and the molar mass |

| |“gram-formula mass”. |Determine the number of moles of a substance, |

| |The mass of one mole of any substance is equal to its molar mass (gram-formula mass). |given its mass and vice versa |

|3. |The percent composition by mass of each element in a compound can be calculated mathematically. |Calculate the percent composition of any element |

| | |in a given compound |

| | |Calculate the percent composition of water in a |

| | |given hydrate |

|4. |Balanced chemical equations show conservation of matter, energy, and charge. |Balance equations, given the formulas for |

| |The coefficients in a balanced equation can be used to determine mole ratios in the reaction. |reactants and products |

| | |Calculate simple mole-mole ratios, given balanced |

| | |equations |

|5. |Types of chemical reactions include synthesis, decomposition, single replacement, and double |Identify the different types of chemical |

| |replacement. |reactions, given their chemical equations |

|Physical Behavior of Matter |

| |Knowledge |Application |

|1. |Physical properties of substances can be explained in terms of chemical bonds and |Predict relative melting and boiling points of compounds|

| |intermolecular forces. |given information about its chemical bonds or strength |

| |Intermolecular forces created by an unequal distribution of charge result in varying |of intermolecular forces |

| |degrees of attraction between molecules. Hydrogen bonding is an example of a strong |Predict relative strength of intermolecular forces of a |

| |intermolecular force that occurs in compounds containing H bonded to F, O, or N atoms. |compound given its melting and boiling points |

| |Physical properties include malleability, solubility, hardness, melting/freezing point, | |

| |and boiling point. | |

|VII. Physical Behavior of Matter – solids, liquids, and gases |

| |Knowledge |Application |

|2. |The three phases of matter (solids, liquids, and gases) have different properties. |Draw and interpret a particle diagram to differentiate |

| |The structure and arrangement of particles and their interactions determine the physical|among solids, liquids, and gases |

| |state (s, l, or g) of a substance at a given temperature and pressure. |Explain phase change in terms of the changes in energy |

| | |and intermolecular distances |

|3. |Heat is a transfer of energy (thermal energy) from a body of higher temperature to a |Distinguish between heat energy and temperature in terms|

| |body of lower temperature. Heat (thermal energy) is associated with the random motion |of molecular motion and amount of matter |

| |of atoms and molecules. |Convert between degrees Celsius and degrees Kelvin |

| |(Heat flows from HOT materials to COLD materials) |(Table T) |

| |Temperature is a measure of the average kinetic energy of the particles in a sample of | |

| |matter. Temperature is NOT energy – it is a measure of heat energy. | |

|4. |The concepts of kinetic and potential energy can be used to explain physical processes |Identify areas of heating and cooling curves that show |

| |that include: fusion (melting), solidification (freezing), vaporization (boiling, |changes in kinetic and potential energy, heat of |

| |evaporation), condensation, sublimation, and deposition. |vaporization, heat of fusion, and phase changes |

| |The kinetic and potential energy changes involved in these physical processes can be |Calculate the heat involved in a phase or temperature |

| |illustrated in a heating curve or cooling curve. |change of a sample of matter using Tables B & T and/or a|

| | |given heating or cooling curve |

|5. |Physical processes like phase changes can be exothermic or endothermic. |Distinguish between endothermic and exothermic phase |

| | |changes |

|6. |Entropy is a measure of the randomness or disorder of a system. A system with greater |Compare the entropy of different phases of matter |

| |disorder has more entropy. | |

|7. |The concept of an ideal gas is a model to explain behavior of gases. |Given a choice of pressure and temperature conditions, |

| |A real gas is most like an ideal gas when the real gas is at low pressure and high |identify those under which gases behave most ideally |

| |temperature. |and/or least ideally |

|8. |The Kinetic Molecular Theory (KMT) states that, for an IDEAL gas, all gas particles |

| |are in random, constant, straight-line motion |

| |are separated by great distances relative to their size (have negligible volume) |

| |have no attractive forces between them |

| |have collisions that may result in a transfer of energy between them, but the total energy of the system remains constant |

|9. |The Kinetic Molecular Theory (KMT) describes the relationships of pressure, volume, |Explain the gas laws in terms of KMT. |

| |temperature, velocity, frequency, and force of collisions among gas molecules. |Solve problems using the combined gas law (Table T) |

| | |Recognize and draw graphs showing P vs. T, V vs. T, and |

| | |P vs. V |

|10. |Equal volumes of gases at the same temperature and pressure contain an equal number of particles. |

|VII. Physical Behavior of Matter – aqueous solutions |

| |Knowledge |Application |

|11. |Physical processes, such as a compound dissolving in a solution, can be exothermic or |Interpret ∆H values for physical processes given in |

| |endothermic. |Table I |

|12. |A solution is a homogeneous mixture of a solute dissolved in a solvent. |Identify the solute and the solvent in a given solution |

| | |Give examples of different types of solutions |

| |The solubility of a solute in a given amount of solvent is dependent on the temperature,|Predict the effect of temperature, pressure, and nature |

| |the pressure, and the chemical natures of the solute and solvent. |of solvent on solubility for a given solute |

| |General rules: |Use a solubility curve to distinguish among unsaturated,|

| |solubility of a solid increases as temperature increases (direct relationship) |saturated, and supersaturated solutions |

| |solubility of a gas decreases as temperature increases (inverse relationship) |Calculate the amount of a specific solute dissolved at |

| |solubility of a gas increases as pressure increases (direct) |different temperatures using Table G |

| |“like dissolves like” – polar solvents dissolve polar solutes; nonpolar solvents | |

| |dissolve nonpolar solutes | |

|13. |Many chemical reactions happen in solution. When different ionic compounds are mixed |Use Table F (Solubility Guidelines) to determine a |

| |together in the same solution, a double replacement reaction may occur and a stable |compound’s solubility |

| |precipitate (insoluble/solid compound) may form. |Determine if a precipitate will form when ionic |

| | |compounds are mixed in solution |

| | |Write and balance chemical equations for double |

| | |replacement reactions |

|14. |The concentration of a solution may be expressed as: molarity (M), percent by volume |Calculate solution concentrations in molarity (M), |

| |(%v/v), percent by mass (%m/v), or parts per million (ppm). |percent by volume, percent by mass, or parts per million|

| | |(ppm) |

| | |Describe how you would prepare a solution from scratch, |

| | |given the desired molarity |

| | |Describe how you would dilute a solution of known |

| | |concentration (must use the equation M1V1 = M2V2) |

|15. |The addition of a nonvolatile solute to a solvent causes the boiling point of the |Compare the freezing and boiling points of solutions of |

| |solution to increase and the freezing point of the solution to decrease. |different concentration |

| |The greater the concentration of solute particles, the greater the increase in b.p. and | |

| |decrease in f.p. | |

|Kinetics and Equilibrium |

|Knowledge |Application |

|1. |The Collision Theory states that a chemical reaction is most likely to occur if reactant particles collide with the proper energy and orientation. |

|2. |The rate (speed) of a chemical reaction depends on several factors: |Use the Collision Theory to explain how factors such as temperature, |

| |temperature, concentration, nature of reactants, surface area, and the |surface area, and concentration influence the rate of reaction |

| |presence of a catalyst. |Ex: Increasing the temperature, surface area, or concentration all |

| |Ionic compounds generally react faster than covalent (molecular) compounds |lead to an increase in the rate of a reaction because they all increase|

| |A catalyst provides an alternate reaction pathway, which has lower |the number of effective collisions between reactant particles. |

| |activation energy than an uncatalyzed reaction. |Explain, in terms of the number of bonds broken, why ionic compounds |

| | |generally react faster than covalent compounds |

| | |Explain how a catalyst speeds up a reaction |

|3. |Energy released or absorbed during a chemical reaction can be represented |Read and interpret a potential energy diagram |

| |by a potential energy diagram. |Draw and label the following parts of a potential energy diagram for |

| |The difference in PE of the products and reactants is called the heat of |both an endothermic and exothermic reaction |

| |reaction ((H) |( PE of reactants and PE of products |

| |(H = PE products – PE reactants |( heat of reaction ((H) |

| |(H values for many chemical reactions are listed in Table I |( activation energy (for both the forward |

| | |and reverse reactions) |

| | |( activation energy with a catalyst present |

|4. |Chemical and physical changes can reach equilibrium |Distinguish between examples of physical equilibria and chemical |

| |Saturated solutions are examples of systems in physical equilibria (aq ↔ s)|equilibria |

|5. |At equilibrium, the rate of the forward reaction equals the rate of the |Describe what is happening to the concentrations or amounts of |

| |reverse reaction and the measurable quantities of reactants and products |reactants and products in a system at equilibrium |

| |remain constant at equilibrium*CARE* |Describe the rates of opposing reactions in a system at equilibrium |

|6. |LeChatelier’s principle can be used to predict the effect of a stress (such|Describe, in terms of LeChatelier’s principle, the effects of stress on|

| |as a change in pressure, volume, concentration, or temperature) on a system|a given system at equilibrium, including: |

| |at equilibrium. |( Changing the temperature/heating/cooling |

| |According to LeChatelier’s principle, a system at equilibrium will “shift” |( Changing the concentration of a reactant or product |

| |to reduce the effects of a stress placed on the system. It will “shift” |( Changing the pressure or volume (this affects systems involving |

| |AWAY from an INCREASE and will “shift”toward a decrease in concentration or|gases) |

| |temperature (“shift” means that either the forward or the reverse reaction |Also be able to explain why any shifting occurs in terms of Collision |

| |will be “favored” (go faster) until the rates are again equal and |Theory |

| |equilibrium is re-established). | |

| |Changing the pressure or volume only affects systems that contain gases | |

|7. |Systems in nature tend to undergo changes toward lower energy and higher entropy. |

|Organic Chemistry |

| |Knowledge |Application |

|1. |Organic compounds contain carbon atoms which bond to one another in chains, rings, and networks to form a variety of structures. |

| |Organic compounds are named using specific IUPAC rules and can be represented using molecular formulas, structural formulas, or condensed |

| |structural formulas. |

|2. |Hydrocarbons are organic compounds that contain only carbon and hydrogen. |Draw structural formulas for alkanes, alkenes, alkynes, given their |

| |*The C–H bonds are considered to be nonpolar covalent bonds.* |IUPAC names |

| |Saturated hydrocarbons contain only single C–C bonds. |Name a hydrocarbon (IUPAC rules), given its molecular formula, |

| |Unsaturated hydrocarbons contain at least one double or triple |structural formula, or condensed structural formula |

| |carbon-carbon bond. |Identify whether a hydrocarbon is saturated or unsaturated, given its |

| | |IUPAC name, structural formula, condensed structural formula, or |

| | |general formula (see Table Q) |

|3. |In a multiple covalent bond, more than one pair of electrons is shared |Determine the TOTAL number of electrons shared in a covalent bond |

| |between two atoms. In a structural formula, each line represents TWO |Determine the number of electron PAIRS (lines) shared in a covalent |

| |shared electrons |bond |

|4. |A functional group is a group of atoms attached to an organic compound that|Identify different kinds of functional groups |

| |gives distinct physical and chemical properties to organic compounds having|Classify an organic compound based on its structural formula, condensed|

| |that group attached to it. |structural formula, or IUPAC name |

| |Organic acids, alcohols, esters, aldehydes, ketones, ethers, halides, |Draw a structural formula with the functional group(s) on a straight |

| |amines, amides, and amino acids are types of organic compounds that differ |chain hydrocarbon backbone, given the correct IUPAC name for the |

| |in the type of functional group they have. |compound |

| |Compounds that have the same functional group have similar physical and |Name any of these organic compounds, given their structural or |

| |chemical properties |condensed structural formulas (see Table R) |

| |Ex: all esters have pleasant odors, all organic acids donate H+ ions in | |

| |solution, all alcohols have low boiling points, etc. | |

|5. |Isomers are organic compounds that have the same molecular formula, but |Determine if two compounds are isomers, given their molecular formulas,|

| |different structures and properties. |structural formulas, condensed structural formulas, or names |

|6. |Types of organic reactions include: polymerization, substitution, |Identify types of organic reactions, given balanced chemical equations |

| |fermentation, addition, combustion, esterification, and saponification. |Determine missing reactants or products in a balanced equation, given |

| |*P.S. FACES* |the type of reaction. |

|Oxidation-Reduction Reactions |

| |Knowledge |Application |

|1. |An oxidation-reduction (redox) reaction involves the transfer of electrons |Determine the number of moles of electrons lost or gained in a redox |

| |from one species (element or ion) to another. |reaction, given the other value |

| |The number of electrons lost equals the number of electrons gained | |

| |(conservation of charge) | |

|2. |Oxidation numbers (states) can be assigned to atoms and ions. |Assign oxidation states to atoms and ions |

| |Changes in oxidation numbers indicate that a redox reaction has occurred. |Determine if a reaction is a redox reaction given the reaction |

| | |equation (Hint: any reaction in which an element is alone (uncombined|

| | |with another element) on one side, but in a compound on the other side|

| | |– it’s a redox reaction!) |

|3. |Losing Electrons is Oxidation (LEO) |Determine which species undergoes reduction (oxidation state goes |

| |Gaining Electrons is Reduction (GER) |down) and which species undergoes oxidation (oxidation state goes up) |

| |Oxidized and reduced species are ALWAYS on the LEFT (reactants) side of the | |

| |equation. | |

|4. |An oxidation half-reaction shows which species is oxidized and the number of|Determine if a given half-reaction is showing oxidation or reduction |

| |electrons it loses (electrons go on the right side of the arrow) |Write and balance oxidation and reduction half-reactions (*remember |

| |A reduction half-reaction shows which species is reduced and the number of |conservation of mass and charge – multiply one or both of the |

| |electrons it gains (electrons go on the left side of the arrow) |half-reactions to make electrons lost = electrons gained if they are |

| | |not equal at first) |

|5. |An electrochemical cell can either be a voltaic cell (a battery) or an |Explain, in terms of atoms and ions, why the anode loses mass and the |

| |electrolytic cell. |cathode gains mass |

| |In both voltaic and electrolytic cells |Compare and contrast voltaic cells with electrolytic cells |

| |( oxidation occurs at the anode (An Ox) | |

| |( reduction occurs at the cathode (Red Cat) | |

| |( the anode loses mass | |

| |( the cathode gains mass | |

|6. |A voltaic cell spontaneously converts chemical energy into electrical |Given a diagram of a voltaic cell, identify and label the cathode, |

| |energy. |anode, salt bridge, and the direction of electron flow |

| |The purpose of the salt bridge is to allow for the migration of ions between|Explain the function of the salt bridge and the direction of positive |

| |half-cells |and negative ion migration |

| | |Write balanced oxidation and reduction half-reactions |

|7. |An electrolytic cell requires electrical energy to produce a chemical |Given a diagram of an electrolytic cell, identify and label the |

| |change. Electrolytic cells can be used for electrolysis (splitting a |cathode, anode, and direction of electron flow. |

| |compound into its elements) and for electroplating (coating something with a| |

| |metal). | |

|Acids, Bases, and Salts |

| |Knowledge |Application |

|1. | |Know the definitions of Arrhenius acids and bases |

| |The behavior of many acids and bases can be explained by the Arrhenius | |

| |Theory. |If given the properties, chemical formula, or name, identify a substance|

| | |as an Arrhenius acid or Arrhenius base. (Use Tables K, L, and T to help|

| |Arrhenius acids produce H+ (hydrogen ions) as the only positive ions in |you remember these definitions. Arrhenius acids begin with H, Arrhenius|

| |aqueous solution. The hydrogen ion may also be written as H3O+ and called|bases are metals + hydroxide ion(s). *Don’t be fooled by alcohols, |

| |the hydronium ion. |which also end in OH, but contain covalent bonds and do not ionize like |

| | |bases do in solution. ALCOHOLS ARE NOT BASES!) |

| |Arrhenius bases produce OH– (hydroxide ions) as the only negative ion is | |

| |aqueous solution. | |

| |(Table E) | |

|2. |Arrhenius acids, Arrhenius bases, and salts (ionic compounds) are all | |

| |electrolytes. An electrolyte is a substance which, when dissolved in |Given names or chemical formulas, identify acids, bases, and salts as |

| |water, forms a solution capable of conducting an electric current |being electrolytes |

| |(electricity). Electrolytes can conduct electricity because they ionize | |

| |(break apart into ions) in a solution. |Determine the relative strength (strong or weak) of an electrolyte given|

| | |information on its ability to ionize in solution. |

| |The ability of a solution to conduct an electric current depends on the |(Strong acids and strong bases are strong electrolytes – Tables K and L |

| |concentration of the ions in it (more ions, more conduction). |list acids and bases in order from strongest to weakest. |

| | |If a salt is soluble, it is a strong electrolyte – Table F can be used |

| | |to determine the solubility of different salts.) |

|3. |In the process of neutralization, an Arrhenius acid and an Arrhenius base | |

| |react to form a salt and water. |Recognize neutralization reactions when given the reaction equation |

| |Acid + Base ( Salt + Water | |

| | |Write neutralization reactions when given the reactants. (Remember that|

| | |this is a double replacement reaction. Just switch the positive ions, |

| | |look up their charges and cross down the subscripts if needed. Then |

| | |balance the equation.) |

|4. | |Calculate the concentration or volume of a solution, using titration |

| |Titration is a laboratory process in which a volume of solution with a |data using the equation |

| |known concentration is added to another solution of unknown concentration.|Ma x Va = Mb x Vb |

| |Titrations are done to determine the concentration of the unknown |(This equation is on Table T) |

| |solution. | |

|5. | |Give the alternate definitions of acids and bases |

| |There are alternate acid-base theories. One such theory states that the |Use this definition to explain why ammonia is considered a base |

| |acid is a proton donor (H+ donor) and the base is a proton acceptor. | |

|6. | |Identify a solution as acidic, basic (alkaline), or neutral based upon |

| |The acidity or alkalinity of an aqueous solution can be measured using the|the pH value OR the relative concentrations of H+/H3O+ and OH– |

| |pH scale. |Describe acidic, basic, and neutral solutions in terms of pH value and |

| |The pH scale measures the concentration of H+/H3O+ in a solution. |relative H+/H3O+ and OH– concentrations |

| |[H+] = 10–pH |Differentiate between strong acids/bases and weak acids/bases given pH |

| |A pH of 1 means that the [H+] = 10–1 = 0.1M |values or ion concentrations |

| |A pH of 3 means that the [H+] = 10–3 = 0.001M | |

| |( Acids have pH values between 0 and 7 | |

| |(the stronger the acid, the lower the pH and the more | |

| |H+) | |

| |[H+] ( [OH–] | |

| |( Neutral solutions have a pH of 7 | |

| |[H+] = [OH–] | |

| |( Bases have pH values between 7 and 14 | |

| |(the stronger the base, the higher the pH and the | |

| |more OH–) | |

| |[H+] ( [OH–] | |

|7. |The pH scale is a logarithmic scale, which means that a change of one pH |Determine the new pH value of a solution given the starting pH and the |

| |unit changes the concentration of H+/H3O+ by a factor of ten |amount of increase or decrease in [H+]/[H3O+] (such as tenfold, a |

| |( tenfold = 10 times = 101 |hundredfold, or a thousandfold) |

| |( hundredfold = 100 times = 102 |Ex: A lake with an initial pH of 6 has been affected by acid rain. The|

| |( thousandfold = 1000 times = 103 |acid rain has caused a hundredfold change in the [H+] concentration of |

| |The exponents represent the CHANGE in pH |the lake. What is the new pH of the lake? |

| |If a solution becomes more acidic, |Answer: pH = 4 |

| |the pH (, and the [H+]/[H3O+] ( | |

| |If a solution becomes more basic, |Determine the amount that the [H+]/[H3O+] would increase or decrease |

| |the pH (, and the [H+]/[H3O+] ( |given a certain change in pH |

|8. |The pH of a solution can be shown by using indicators. |Interpret changes in acid-base indicator color |

| |An indicator is a substance that changes color depending on the |Explain how different indicators can be used to distinguish between |

| |concentration of hydrogen/hydronium ions in a solution. |solutions with different pH values |

| | |Identify appropriate indicators that can be used to show changes in pH |

| | |values, such as during a titration, given starting and ending pH values |

|Nuclear Chemistry |

| |Knowledge |Application |

|1. |The stability of an isotope is based on the ratio of the neutrons and protons in the nucleus. |

| |Usually when the ratio is not 1:1, the nucleus gets a little unstable and starts spitting out particles so that it will have a more stable 1:1 |

| |ratio. |

| |Although most nuclei are stable, some are unstable and spontaneously emit radiation. We call these unstable isotopes radioactive isotopes, |

| |radioisotopes, or nuclides. |

|2. |Spontaneous decay (natural emission of radiation) by a nuclide (radioactive |Determine decay mode and write nuclear equations showing alpha decay,|

| |isotope) involves the release of particles and/or energy from the nucleus. |beta decay, positron emission, and gamma radiation (*Remember to put |

| |Each radioactive isotope has a specific decay mode (the kind of particle or |radioactive emissions on the RIGHT side of the arrow – if something |

| |energy it gives off from its unstable nucleus) (Tables N and O!) |is released, it goes on the right) |

| |( alpha decay: release of alpha particles | |

| |( beta decay: release of beta particles |Compare and contrast the 4 different types of radiation in terms of |

| |( positron emission: release of positrons |mass, charge, ionizing power, and penetrating power. |

| |( gamma radiation : release of gamma rays | |

| |These emissions differ in mass, charge, ionizing power, and penetrating | |

| |power. | |

|3. |Each radioactive isotope has a specific half-life (rate of decay). The |Calculate the initial amount, the fraction remaining, time elapsed, |

| |half-life is the time it takes for half of the radioisotope to |or the half-life of a radioactive isotope, given the other variables |

| |decay/transmutate into something more stable). (Table N) | |

|4. |Nuclear reactions are represented by equations that include symbols for |Complete nuclear equations and predict missing particles in nuclear |

| |elements and radioactive emissions (with mass number in upper left and |equations |

| |charge/atomic number in lower left) |Write nuclear equations given word problems |

| |These reactions show conservation of mass and charge | |

|5. |A change in the nucleus of an atom that changes it from one element to |Distinguish between natural transmutation (one reactant) and |

| |another is called transmutation. This can occur naturally or can be done |artificial transmutation (two reactants) given nuclear equations |

| |artificially by bombarding the nucleus with high-energy particles. | |

|6. |Types of nuclear reactions include fission and fusion. Fission and fusion | |

| |can be natural or artificial transmutations. |Compare and contrast fission and fusion reactions. |

| | |Distinguish between fission and fusion reactions given nuclear |

| | |equations |

|7. | |Compare and contrast chemical reactions and nuclear reactions |

| |Nuclear changes convert matter into energy |Describe benefits of using nuclear fission |

| |(E = mc2) | |

| |Energy released during nuclear reactions is much greater than the energy | |

| |released during chemical reactions. | |

|8. | |Describe the risks and problems associated with using radioactive |

| |There are risks and problems associated with radioactivity and the use of |isotopes |

| |radioactive isotopes, including: biological exposure, long-term storage and | |

| |disposal problems, and nuclear accidents which release radioactive materials | |

| |into the environment. | |

|9. | |

| |In addition to using nuclear fission for nuclear power, radioactive isotopes have other beneficial uses in medicine and industrial chemistry, |

| |including: |

| |( radioactive dating (ages of once-living things can be found from the ratio of C-14 to C-12 in the remains; ages of rocks can be found from the |

| |ratio of U-238 to Pb-206) |

| |( tracing chemical and biological processes (radioactive tracers can be injected into the body and then x-rayed. The radioactive substance will |

| |show up on the x-ray and if there are problems, they can be detected easily) |

| |( detecting and treating of disease (Sr-90: diagnosing and treating bone cancer; I-131: diagnosing and treating thyroid disorders; Co-60: |

| |cancer treatment |

| |( radiation can be used to kill bacteria in foods (used with spices, meats, produce) |

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