Chapter 19 Notes - Physical Science with Mrs. Furstenberg



Parts of the Atom

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PERIODIC TABLE BASICS

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Neutral Atoms and Ions

If the number of ____________ equals the number of ___________ the atom is _____________.

If not, it is an ________________.

|Protons |Electrons |Neutral or Ion? |Charge |

|6 |6 | | |

|8 |7 | | |

|3 |4 | | |

|2 |2 | | |

|1 |1 | | |

|9 |10 | | |

Isotopes

An _______________ is a variation of an element. It has the ____________ number of ___________ (same element), but a different number of _______________ (different isotope).

Isotopes of the same element have different __________________.

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History of the Periodic Table

How Should Elements be arranged?

John Dalton attempted to classify elements on the basis of their ________ ___________.

Then, in 1829, __________________ __________________ was the first to classify elements into groups based on John Dalton’s assertions.

He groups the elements with similar ____________ ___________ into clusters of three called ___________.

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Mendeleev’s Periodic Table

Arranged in order of _____________ ______________ ______________.

❖ Didn’t work! Why? Mass doesn’t always go up. Ex: Co, Ni

THEN, he grouped elements together with similar ____________.

Modern Periodic Table:

*Increasing _______ _________

1. Rows are called __________.

a. Cycles from left to right with properties

b. Represents a shell (n). (Think energy levels!)

2. Columns are called _________.

a. Similar properties!-reactivity groups

b. Same number of valence electrons

Periodic Table Groups

Alkali Metals

✓ Group 1

✓ Hydrogen is not a member, it is a non-metal.

✓ ___ electron in the outer shell.

✓ Soft and silvery _______.

✓ Conduct ___________.

Alkaline Earth Metals

✓ Group 2

✓ ___ electrons in the outer shell.

✓ Conduct electricity.

Transition Metals

✓ Groups in the middle.

✓ Good conductors of heat and electricity.

Boron Metals

✓ Group 3

✓ _______ electrons in the outer shell.

✓ Most are metals.

Carbon Family

✓ _____ electrons in the outer shell.

✓ Contains, metals, metalloids and a non-metal (Carbon)

Nitrogen Family

✓ Group 5

✓ ___ electrons in the outer shell.

✓ Can __________ electrons to form compounds.

✓ Contains metals, metalloids, and non-metals.

Oxygen Family

✓ Group 6

✓ ____ electrons in the outer shell.

✓ Contains metals, metalloids, and non-metals.

✓ Reactive

Halogens

✓ Group 7

✓ ____ electrons in the outer shell.

✓ All are ________________.

✓ Very reactive, often bond with elements from __________.

Noble Gases

✓ Group 8

✓ Non-metals

✓ Exist as gases.

✓ ______ electrons in the outer shell = FULL

✓ Helium has only 2 electrons in the outer shell = FULL.

✓ NOT REACTIVE.

|Element |Period |Group |Family |# Electrons in Outside Shell |

|Silicon | | | | |

|Sulfur | | | | |

|Krypton | | | | |

|Aluminum | | | | |

|Sodium | | | | |

|Calcium | | | | |

Drawing Bohr Diagrams

Named after _______ __________.

Bohr diagrams allow us to see how many ___________an atom has.

Step 1: Write the element symbol and circle it. This is your nucleus.

Step 2: Draw the electron shells. (This is the period #)

Step 3: Draw the electrons in each shell.

Remember: 1st shell holds a maximum of 2 electrons.

2nd shell holds a maximum of 8 electrons.

3rd shell holds a maximum of 18 electrons.

4th shell holds a maximum of 32 electrons.

Example:

Carbon

You Try it:

1. Lithium 2. Aluminum 3. Neon

Chemical Bonds

Valence Electrons

o Contained in the outermost region of the electron cloud (called the valence shell)

o The electrons available to be __________, ___________ or ___________ in the formation of chemical compounds.

□ Group 1- one valence electron

□ Group 2- two valence electrons

□ Groups 13-18- Have the number equal to the group number minus____.

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Octet Rule: Atoms form chemical bonds with other atoms by either ________, __________ or ___________ electrons in order to have a complete set of _______ valence electrons.

□ Exceptions: the 1st energy level only needs ___ electrons.

(Stable Atoms have FULL or EMPTY valence shells!)

Complete the table: Elements, number of valence electrons, and number needed to complete the octet

|Element |# of Valence electrons |Number of valence electrons needed to complete |

| | |the octet. |

|H |1 | |

|He |2 | |

|Li | | |

|Be | | |

|B | | |

|C | | |

|N | | |

|O | | |

|F | | |

|Ne | | |

|Ar | | |

LEWIS STRUCTURES/LEWIS DOT DIAGRAMS

Showing the element symbol surrounded by 1 to 8 dots representing the valence electrons:

• Atoms are most stable with 2 or 8 valence electrons.

• Noble gases already have 8 electrons, so they usually don’t form bonds with other atoms.

• Other atoms form bonds so that they can share electrons to reach that stable number of _____ electrons.

1. Write the element symbol.

2. Determine the # of valence electrons the element has. (Think Column #)

3. Place electrons around the atom (singly and then pairing them up)

SINGLE ATOM EXAMPLES:

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Boron: Neon:

Fluorine: Oxygen:

We try it!

Draw the Lewis Dot Diagrams of the following elements:

1. Lithium

2. Silicon

3. Sulfur

4. Helium

Bonding

|COVALENT BONDING | |

| | |

|A covalent bond forms between 2 elements because they | |

|__________ one or more pair of valence electrons | |

|between the two atoms. | |

|This allows both of the atoms to receive a ________ | |

|______ of valence electrons. | |

|The two atoms are more ________ together than when | |

|they are apart. | |

Together as a class we will practice drawing covalent bonds with Lewis Dot Structures.

Then, draw the Lewis Structures for the following covalent compounds:

1. H2O

2. NH3

3. HF

Ionic Bonds

Formed when ions are __________ from one atom to another.

Ions-atoms that either _____ or ______ an electron. Atoms that accept electrons will have a _________ charge. Atoms that donate electrons will have a _________ charge.

-Ionic vs. Covalent: depends on how close an atom is to eight valence electrons. ________________ with one valence electron have high tendency to give up an electron. ____________ with seven have a high tendency to take an electron. Elements with ___ to ____ electrons tend to _______ electrons by forming covalent bonds.

A ____________________occurs between _____________________________.

An _______________occurs between a __________ and a ________________.

Drawing Lewis Structures for Ionic Compounds

1. First draw the Lewis Dot Diagram for EACH atom individually.

2. Determine which atom will donate an electron.

3. Draw an arrow to show where the electron will be transferred to.

Examples:

1. Sodium + Fluorine

2. 2 Lithium + Oxygen

TRY IT ON YOUR OWN:

1. Lithium + Fluorine

2. Hydrogen + Chlorine

3. 2 Hydrogen + Oxygen

| |

|Oxidation Numbers |

|Oxidation Numbers- indicates how many _________ are lost, gained, or shared when bonding occurs. |

|If an atom has a tendency to _____ an electron away, it is _________. |

|It is GOOD (+) to GIVE. |

|Known as a __________. |

| |

|If an atom has a tendency to ________ an electron away from another atom, it is ___________. |

|It is BAD (-) to TAKE. |

|Known as an __________. |

| |

|Sodium (Na) always ionizes to become Na+, so its oxidation number is _____. |

|Oxidation numbers and the periodic table: [pic] |

|Complete the chart: |

|Atom |

|Electrons gained or lost |

|Oxidation number |

| |

|K |

|Loses 1 |

|1+ |

| |

|Mg |

|Loses 2 |

|2+ |

| |

|Al |

| |

| |

| |

|P |

| |

| |

| |

|Se |

| |

| |

| |

|Br |

| |

| |

| |

|Ar |

| |

| |

| |

Writing Chemical Formulas

✓ Formed when atoms combine in certain ________ with other atoms.

✓ Ionic Compoud-compounds made of atoms that form __________ __________.

o Atoms ________ and ________ electrons.

✓ Example: Na+ + Cl- NaCl

o All compounds have a charge of _________; they are neutral.

| Writing | |The ionic charges combine in the simplest ratio possible that allows the positive |

|Ionic |There are three rules to writing a correct ionic formula: |and negative charges to balance out (____________________). Add the subscripts |

|Compounds | |multiplied by the oxidation numbers and make sure it equals zero. |

| | |The cation (positive ion) is _____________________________!! |

| | |Negative ion or anion is _____________________________!!! |

| |Steps for Writing Ionic Formulas: |Example: |

| |Find the oxidation numbers of each element of the compound. If it|Write the formula for a compound that is made of iron (III) and oxygen. |

| |is a transition metal, the roman numeral (III) tells you the | |

| |oxidation number. Ex. Iron (3+). | |

| |Determine the ratios of each element so that the compound is | |

| |neutral (no charge) and write the chemical formula. | |

| | | |

| | | |

| | | |

| | | |

| | | |

| | | |

| | | |

| | | |

| |Examples using the Secret Weapon: Criss-Cross Rule |

| |Examples: (Make sure to determine oxidation numbers) |

| | |

| |sodium + phosphorus ( |

| | |

| |lithium+ chlorine ( |

| | |

| |aluminum + fluorine ( |

| | |

| |tin (IV) + oxygen ( |

| | You Try it! |

| |Write the chemical formula for each compound: |

| |Lithium + Oxygen→ |

| |Calcium + Oxygen→ |

| |Barium + Bromine→ |

| |Potassium + Chlorine→ |

| |Potassium + Sulfur→ |

| |Sodium + Iodine→ |

| |Strontium + Bromine→ |

| |Ionic compounds containing polyatomic ions |

| |Write the chemical formula and oxidation of the positive ion. If the positive ion is monoatomic, use the periodic table to find oxidation number. If|

| |the positive ion is polyatomic, use the chart below. |

| |Write the chemical formula and oxidation number for the negative ion. Again, use periodic table of monoatomic and chart if polyatomic. |

| |Add the oxidation numbers and make sure they equal zero. Figure out how many of each ion you will need so that the compound has no charge. You can use|

| |the criss-cross rule. |

| |Examples-write the chemical formula: | |

| |Aluminum + sulfate→ | |

| |Magnesium + carbonate→ | |

| |Ammonium+ hydroxide→ |Tip: The positive ion is always written FIRST! It should |

| |Hydronium + fluorine → |also be first in the formula. |

Polyatomic ions:

|Oxidation Number |Name |Formula |

|1+ |Ammonium |NH4+ |

|1- |Acetate |C2H3O2- |

|2- |Carbonate |CO32- |

|2- |Chromate |CrO42- |

|1- |Hydrogen carbonate |HCO3- |

|1+ |Hydronium |H3O+ |

|1- |Hydroxide |OH- |

|1- |Nitrate |NO3- |

|2- |Peroxide |O22- |

|3- |Phosphate |PO43- |

|2- |Sulfate |SO42- |

|2- |Sulfite |SO32- |

Naming Compounds

Metal + Non-Metal = Ionic Compound

| |Naming Ionic Compounds can be broken down into two scenarios: |

|Naming |Scenario #1: Naming ionic compounds with monoatomic ions |

|Ionic | |

|Compound| |

|s | |

| |Write the name of the first |Write the root name of the second element. For example (Cl), chlor- is the root. |

| |element in the compound. THEN, LEAVE IT ALONE!!! |Simply subtract the –ine and ADD –ide. Chlor- |

| | |becomes chloride. |

| |Examples: |ELEMENT |ION NAME |

| |KCl | | |

| | | | |

| |AlP | | |

| | | | |

| |Ca3N2 | | |

| | | | |

| |MgBr2 | | |

| | |Sulfur |Sulfide |

| | |Nitrogen |Nitride |

| | |Phosphorus |Phosphide |

| | |Fluorine |Fluoride |

| | |Oxygen |Oxide |

| | |Chlorine |Chloride |

| | |Bromine |Bromide |

| | |Iodine |Iodide |

| | |Selenium |Selenide |

| |Scenario #2: Naming ionic compounds with polyatomic ions |

| |Write the name of the positive ion first. Use the periodic table or the polyatomic ion chart to find its name. |

| | |

| |Write the name of the negative ion second. Use the periodic table or the polyatomic ion chart to find its name. Add –ide if the negative ion is |

| |monoatomic. |

| | |

| |Examples: |

| |Mg(CO3) |

| | |

| |(NH4)F |

| | |

| |Ca(NO3)2 |

| | |

| |(H3O)(OH) |

| | |

Naming Covalent Compounds

Non-Metal + Non-Metal = Covalent

|Naming and Writing Covalent Compounds |

|Binary compounds-consists of only ____ elements |

|Naming &| |Mono- |1 |

|Writing |When we name covalent compounds we will use prefixes. The subscripts beside each atom will be converted to a prefix | | |

|Covalent|tacked onto the front of each element name (except the first element named never uses the prefix __________). | | |

|Compound| | | |

|s |The second element receives and prefix appropriate and ends with the suffix _________ just like in naming ionic | | |

| |compounds. | | |

| | | | |

| | |Di- |2 |

| | |Tri- |3 |

| | |Tetra- |4 |

| | |Penta- |5 |

| | |Hexa- |6 |

| | |Hepta- |7 |

| | |Octa- |8 |

| | |Nona- |9 |

| | |Deca- |10 |

| |Practice Naming Formulas |Practice Writing Formulas |

| |CO2 |Hexacarbon hexahydride |

| |P3O5 |Boron trifluoride |

| |C4H9 |Pentaiodide decahydride |

| |H2O |Carbon monoxide |

Formula Mass/Molar Mass

Formula mass (atomic mass unit, amu)

• the sum of the average atomic masses of all atoms represented in its formula (i.e. Carbon= 12.01 amu)

• A way to compare the masses of different compounds.

Example:

1. H2SO4

2. Li2O

Find the formula mass for each of the following compounds:

|Compound |Formula Mass |

|1. NaCl | |

|2. H2O | |

|3. CO2 | |

|4. HCl | |

|5. K2S | |

|6. MgCl2 | |

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Tells us the sum of _______ +________

Copper

Tells us the # of _______ or _________

Steps for creating multiple atom Lewis Dot Diagrams:

1. First, draw Lewis Dot Diagrams for EACH atom individually.

2. Determine where bonds can form between the atoms. (Circle These!)

3. Arrange the atoms for the bonds, and draw the combined Lewis Dot structure.

4. Replace the shared electrons with single bonds.

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