Chapter 1 Fun with the Periodic Table



INTRODUCTION TO CHEMISTRY

JULY 2008

RYAN McGRATH

JOHN O’DONNELL

ALISA PRUDENTI

REVISED 2014

SCOPE

Introduction to Chemistry

1. Fundamentals of Chemistry

A. Defining matter

1. Differentiating matter and energy

2. Phases of matter

3. Composition: Substances and mixtures

a. Elements / Compounds, Homogenous/Heterogeneous

b. Separation of a Mixture

4. Chemical and physical properties

5. Chemical and physical changes

6. Conservation of mass

B. Energy

1. Forms of energy

2. Endothermic and exothermic reaction

3. Mass and energy are conserved in a chemical reaction

4. Temperature:

a. average kinetic energy of particles

b. Kelvin temperature = oC + 273

2. Atomic Structure

A. Structure of the atom: particles and the nucleus

1. Comparison of proton, neutron and electron

a. mass

b. charge

c. location

d. symbol

2. The atomic nucleus

a. determining the number of protons and neutrons

b. definition of isotope

c. Qualitative Atomic Mass (which isotope is more abundant)

B. Electronic structure

1. Energy levels and the Periodic Table

2. Valence electrons

3. Atomic spectra

a. ground state vs. excited (identify)

b. explanation of energy in electron transitions (absorb/emit)

c. qualitative (color)

3. The Periodic Table

A. Organization of the Periodic Table (using the Regents Periodic Table)

1. Information in each box of the periodic table

a. symbol

b. atomic number

c. atomic mass (average isotopic)

d. charge of ions formed

e. principal energy level valence electron configuration

1. valence electron configuration changes for each element in a period

2. valence electron configuration is the same for each element in a group

2. Periodic trends: horizontal rows (periods)

a. metals ( metalloids ( nonmetals (decrease in metallic properties)

b. use table S for trends

3. Periodic trends: descending vertical columns

a. increase in metallic properties

b. use table S for trends

4. Properties of metals and nonmetals

a. metals

1. conduct electricity

2. conduct heat

3. have luster and gray color

4. are malleable and ductile

b. transition metals

1. location

2. ions produce colored solutions

c. metalloids

1. found on either side of the “stairsteps” (except Aluminum).

2. characteristics of both metals and nonmetals

d. nonmetals

1. are nonconductors of heat and electricity

2. are brittle as solids

B. Groups (families) of elements

1. Halogens (group 17)

a. very reactive nonmetals

2. Noble gases (group 18)

a. complete valence electron levels

b. generally unreactive (although Ar, Kr and Xe can be made to form compounds)

3. Alkali metals (group 1)

a. most reactive metals

4. Alkaline earth metals (group 2)

a. very reactive metals

4. Chemical Formulas and Equations

A. Significance of the formula

1. identification of elements present (qualitative)

2. ratio of atoms of each element present (quantitative)

a. subscript represents number of atoms of element to its left

b. ratio of subscripts is the ratio of each atom present

3. the sum of oxidation numbers of any compound is 0

4. subscripts are added to each element when necessary to reach sum of 0

5. subscripts are reduced to lowest terms

6. Naming of simple ionic compounds (no stock system)

a. Name ( Formula, Formula ( Name

B. Polyatomic ions in chemical formulas (Table E)

1. atoms that form stable ionic groups

C. Balancing equations

1. Reactant(s) ( product(s)

2. mass is conserved in all chemical reactions

a. equations represent conservation of mass

b. coefficients represent number of atoms or molecules present

c. all atoms present in reactants are present in products

d. determine coefficients in an unbalanced equation

3. Equation types

a. synthesis

b. decomposition

c. single replacement

d. double replacement

e. combustion

5. Mathematics of Stoichiometry

A. Formulas and moles

1. Chemical formulas represent

a. one atom, molecule or formula mass of that substance in atomic units

2. Empirical and molecular formulas

a. empirical formula is lowest whole number ratio of atoms of each element

b. it is calculated from masses of substances present in laboratory experiments

c. gram formula mass calculation

d. convert mass to moles

e. mole conversions

f empirical ( molecular

6. Bonding

A. Bonds form, energy released (absorbed when breaking)

B. Ionic Bonding: the Octet rule

1. Stable octet and noble gases

2. Tendency of atoms to combine to achieve stable octet

a. Metals lose valence electrons to achieve stable octet at lower principal energy

level

b. Losing electrons, metals form positive ions

c. Nonmetals gain valence electrons to achieve stable octet at same principal

energy level

d. Gaining electrons, nonmetals form negative ions

e. Ionic dot diagrams for binary compounds

C. Covalent bonding

1. sharing of electrons between nonmetal atoms

2. octet rule

3. dot diagrams (no multiple bonds)

7. Thermochemistry

A. volume, measured in liters (L) or milliliters (mL)

1. space occupied by gas molecules

Energy

1. Forms of energy

a. Kinetic / Potential

b. Temperature: average kinetic energy of particles

2. Endothermic and exothermic reaction

3. Mass and energy are conserved in a chemical reaction

4. Calculate amount of heat absorbed/released during temperature changes of a substance.

a. definition of Specific Heat capacity

b. q = mC∆T

D. Phase change

1. boiling/condensation : liquid/ gas

a. energy absorbed is used to separate molecules so far that they have no attraction

b. no temperature change since kinetic energy is not increased

c. gain/loss of heat is defined as heat of vaporization

d. q= mass x Hv

2. freezing/melting : solid/liquid

a. energy absorbed breaks orderly arrangement

b. no temperature change since kinetic energy is not increased

c. gain/loss of heat is defined as heat of fusion

d. q= mass x Hf

3. Qualitatively identify phases, phase changes and energy changes on a heating curve.

8.. Solutions

1. homogeneous mixture (the same throughout)

a. may include combinations of phases

b. solute: the substance that dissolves

c. solvent: the substance in which the solute dissolves

2. Temperature (Table G)

a. most solids increase in solubility with increasing temperature

b. all gases decrease in solubility with increasing temperature

3. Solubility curves

a. all solutions are in 100 g (mL) of water

1. concentrations on curve line are saturated

2. concentrations below curve line are unsaturated

3. concentrations above curve line are supersaturated

b. problems involving solutions cooling to form precipitate

1. calculate mass of precipitate formed when solution cooled.

4. Solubility rules (Table F)

a. use guidelines on reference tables

b. insoluble substances form precipitates

5. Solution concentration

a. dilute

1. little solute in a given amount of solvent

2. occurs in two cases

a. adding small amount of a very soluble solute

b. adding any amount of a relatively insoluble solute

b. concentrated

1. a large amount of solute in a given amount of solvent

2. insoluble substances cannot form concentrated solutions

c. molarity

1. moles of solute per liter of solution (just solve for M)

9. Acids and Bases

A. Arrhenius acids and bases

1. Acid is a substance that adds H+ ions as only positive aqueous ion

2. properties of acids

a. sour taste

b. conduct electricity (are electrolytes)

c. change color of indicators

d. react with bases to form salts and water

3. Bases is a substance that adds OH- ions as only negative aqueous ion

4. properties of bases

a. bitter taste

b. slippery feel

c. conduct electricity (are electrolytes)

d. change color of indicators

e. react with acids to form salts and water

B. Neutralization

a. Acids neutralize bases/ bases neutralize Acids to produce salt and H2O

C. pH scale

1. strengths of acids and bases

2. the acidity of an aqueous solution is measured in terms of relative presence of H+ and OH- ions

a. given [H+] , determine pH ([H+] given as 10x molar)

a. determine new pH when [H+] changes by a factor 10,100, 1000

4. indicator

a. determine color in a solution given a specific pH (Table M)

5. salts are formed in a neutralization reaction

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