Chemistry Notes

NAME

CHEMISTRY NOTES

CHAPTER 5

PATTERNS AND COMPOUNDS (P 136 – 177)

Number / Item / Value / Your Mark
On Top / This page with your name on it
1 / Ch 9 Review / 5
2 / The Periodic Table/Bohr-Rutherford Chart / 42
3 / Notes 5.1 Looking for Patterns, 5.2 Forming Compounds, Naming Binary Compounds, Chemical Formulas & Eq’ns / 20
4 / P 145 Lewis Dot Diagrams (in notes) – if done / 10
5 / CYU – P 146 1-6 / 18
6 / CYU – P 154 2-4 / 14
7 / BLM 5-1 / 10
8 / BLM 5-2 / 10
9 / BLM 5-5 / 14
10 / BLM 5-6 / 24
11 / BLM 5-7 / 12
12 / BLM 5-9 / 19
13 / BLM 5-10 / 24
14 / BLM 5-11 / 15
15 / BLM 5-13 / 10
16 / BLM 5-14 / 61
17 / BLM 5-15 / 30
18 / BLM 5-16 / 21
19 / BLM 5-17 / 21
20 / 5-3 PRACTICE SHEET (2 pages) / 50
21 / BLM 5-18 / 17
22 / BLM 5-19 / 26
23 / BLM 5-20 / 16
24 / 5-4 PRACTICE SHEET #1 AND #2 / 37
25 / BLM 5-21 / 16
26 / BLM 5-22 (2 PAGES) / 31
573
Percentage / 100

Gr 9 Chem Review

Atomic Structure

John Dalton’s Atomic Theory (1808)

1. All matter is made of atoms

2. Each element has its own kind of atom with the same atomic mass

3. Compounds (molecules) are atoms of different elements linked together

4. Atoms cannot be created or destroyed

Atoms consist of three different subatomic particles

protons – positively charged, located in the nucleus

neutrons – neutral, located in the nucleus

electrons – negatively charged, located outside the nucleus

Bohr Rutherford Model (copy example on overhead)

Atomic # = number of protons inside the nucleus

Mass # = the total number of protons and neutrons in the nucleus

To calculate

The number of protons = atomic number

The number of electrons = atomic number

The number of neutrons = mass number – atomic number

An isotope is an element that can have varying mass numbers due to differences in the number of neutrons

The periodic table is also placed in vertical columns. These are families or groups. They have similar chemical and physical properties.

Nobel Gases

· do not react under normal conditions

· outer electron shell is full

· He, Ar, Ne, Kr, Xe, Rn

Alkali Metals

· One to on reaction with Hydrogen

· Most active metals

· very abundant on earth

· 1 electron in the outer shell

· Na is the most common

· Li, Na, K, Rb, Cs, Fr

Alkaline Earth Metals

· Quite reactive

· quite abundant on earth

· 2 electrons in outer shell

· Be, Mg, Ca

Halogens

· One to one reaction with Hydrogen

· outer shell requires 1 electron to be complete

· most reactive non metal

· Cl is the most common

· F, Cl, Br, I, At

Hydrogen

· can act as a metal or non metal

· great reactivity

· least dense gas

· flammable

All atoms are neutrally charged because they have the same number of protons and electrons

Organizing the Periodic Table

Initially the elements were organized alphabetically. This was not a useful method because:

1. the elements had nothing in common with the elements that were listed next to them

2. the whole list had to be arranged when a new element was discovered.

Dmitri Mendeleev (1855)

1. arranged the periodic table according to atomic mass and element properties

2. left blank spaces where new (undiscovered) elements would likely fit

At the turn of the century, the periodic table was rearranged according to atomic number.

The periodic table is divided by a step like line. On the left are the metal; on the right are the non metals. The elements near the line are metalloids and contain properties of both metals and non metals.

Metal: shiny, solid, high density, good conductor, malleable, few outershell electrons

Non metal: dull, gaseous, low density, poor conductor, brittle, many outershell electrons

The periodic table is arranged in horizontal rows called periods. The elements in each period have the same number of electron shells.

Main element in organic compounds

Notice as you move from left to right on the periodic table

1. the atomic number increases

2. the atomic mass increases

3. elements change from metal, metalloid, non metal, nobel gas

The Periodic Table

Bohr-Rutherford Model Practice

Name: ______/84

Part A: Complete the Chart

Element
Name / Symbol / Atomic Number / Atomic Mass (2 decimals) / # of Electrons / # of Protons / # of Neutrons
0
Phosphorus
9
Mercury
Ne
Silicon
Oxygen
28
Beryllium
Tungsten
65

/33

How is the number of protons determined?

How is the number of electrons determined?

How is the number of neutrons determined?

/36

Part B: Draw the Bohr-Rutherford diagram and fill in the missing information. Include information below your diagram about how many protons, neutrons, and electrons the atom has.

1. H
P = _____ e- = _____ N = _____ / 2. He
P = _____ e- = _____ N = _____ / 3. Li
P = _____ e- = _____ N = _____
4. Be
P = _____ e- = _____ N = _____ / 5. F
P = _____ e- = _____ N = _____ / 6. Ne
P = _____ e- = _____ N = _____
7. Na
P = _____ e- = _____ N = _____ / 8. Cl
P = _____ e- = _____ N = _____ / 9. Ar
P = _____ e- = _____ N = _____
10. K
P = _____ e- = _____ N = _____ / 11. Ca
P = _____ e- = _____ N = _____ / 12. Cu
P = _____ e- = _____ N = _____

/48

5.1 LOOKING FOR PATTERNS IN CHEMICAL REACTIVITY

Valence Shell – The that an atom has.

Valence Electrons – the found in the atom’s shell.

Chemical Reactions

Atoms – NEVER gain or lose . (if they did we could make gold)

- almost NEVER gain or lose . (OK radioactive elements do)

- Often gain or lose . The number of valence electrons give the element it’s and properties.

Atomic Diagrams (draw Na diagrams below )

Bohr-Rutherford Diagram Lewis Dot Diagram

(Electron Dot Diagram)

Since only the valence e- matter for properties, it becomes a waste of time to draw other electrons. Lewis Dot Diagrams are faster and easier to draw, and they show all the important information.

Drawing Lewis Diagrams:

1) Write the element’s symbol

2) Draw the valence electrons clockwise around the symbol.

Read and complete p. 145 in the space below

IONS

Keep an eye-on those electrons!!

Key Eye-dea: Atoms don’t want valence shells .

Cations – charged atoms that have electron(s).

- (1 valence e-) easily the e-!!

· Very

- (2 valence e-) them if possible.

- Reactivity down the family (group) because of the distance from e- to proton(+). It’s the e-.

- Generally form .

Anions – charged atoms that have electron(s)

- (7 valence e-) electrons forcefully!!

- form anions.

Noble Gases – with other atoms because their valence shells are already .

REMEMBER:

Do pg 146 #1 – 6 and BLM’s 5-1 to 5-7

5.2 Forming Compounds

Atoms vs. Ions

Atoms

- # of = # of

- charge

Ions

- atoms which have or electrons

- have on them

Ex. If Mg loses 2 electrons it becomes

(Remember metaLs Lose electrons)

Ex. If Cl gains 1 electron it becomes

(non-metals gain electrons)

Chemical Bonding

Atoms always bond with each other through the electrons. There are 3 ways in which atoms bond:

1) Ionic bonds – between .

- Metals happily their electrons to non-metals which happily the electrons.

- Electrons transfer making two IONS which stick together.

- are formed.

- Ionic compounds:

o have points. (strong ionic bonds)

o easily in water.

o are – materials that conduct when molten or when dissolved (aqueous).

2) Covalent bonds – between two .

- Electrons are by the non-metals.

- sharing electrons allows each atom to have a for short periods of time.

- are formed.

- molecular compounds:

o have points (weaker bonds)

o don’t conduct

o don’t as easily in water.

- there are 2 types of covalent bonds:

a) covalent – the atoms are elements. Electrons are not shared by the atoms. One atom gets the electrons for a longer time. Ex) H2O

b) covalent – the electrons are because both atoms are the same element. Ex) N2, O2, F2, Cl2, Br2, I2 are the only examples. They are called . See the “7” pattern in the periodic table.

3) bonding – between atoms that are .

- Metals allow their electrons to be from one atom to the next.

- (e-) easily move from atom to atom.

Lewis Diagrams for Bonding:

Ionic Bonding – atoms become ions. Metal donates e- to non-metal.

[Li]+ [:H]-

If the metal wants to donate 2 e-, it may need to find 2 non-metals to accept one each.

[:Cl:]- [Mg]2+ [:Cl:]-

OR find a non-metal that will take 2 e-

[Mg]2+ [:O:]2-

The charge that the ion has is often referred to as the combining capacity of the ion.

Try: HCl,

Covalent Bonding – non-metals share electrons. Line up atoms so that they can fill each others outer shells.

H2O H:O:H O has 6 valence e-, H has 1

Try these: CH4, Br2, O2

Comparing Types of Bonding

Ionic / Covalent / Metallic
Called ionic compounds / Called molecular compounds / Names with the metal name
Metals with non-metals / Non-metals only / Metals only
Metals donate e- to non-metals forming oppositely charged ions which stick / Non-metals share e- so that for at least some of the time they will have 8e- / Metals play “hot potato” with e- so e- can flow freely from one atom to the next.
The overall cation charges must balance with the anion charges.
Ex)
Mg+2 with Cl- means we have
MgCl2 / Two types of sharing
1) non-polar covalent – sharing of e- is fair. Occurs in diatoms N2, O2, F2, Cl2, Br2, H2, I2
2) polar covalent – sharing of e- is unfair. One atom hogs electrons more than the other. Ex) H2O / Good for wires. Add an e- to one end and another will pop out the other end.
Very high melting point / Are often liquids or gases / Solids
MOSTLY dissolve in water. Ions break apart. / Polar solutes dissolve in polar solvents
Non-polar solutes dissolve in non-polar solvents / Don’t tend to dissolve
Conduct electricity when dissolved or when molten (Electrolytes) / Don’t conduct electricity very well. / Conduct electricity very well.

Do P 154 2-4 and BLM 5.8 – 5.11, 5.13

5.3 Naming Binary Ionic Compounds

Binary – having only of atoms

Ionic Compounds – electrons are from to . The ions stick together. (pg 148-149 diagrams)

Reading the Name from the formula:

1) Binary compounds usually end it “ ”.

2) Write first and second adding “ ” to the .

Examples: CaCl2 is called

MgBr2 is called

Writing the formula from the name:

1) Determine that the compound is .

a. (metal and a non-metal.) Eg. Calcium & Chlorine

2) Temporarily mark the combining capacity ( ) on each atom. Eg. Ca2+ Cl-

a. Note: in case of a charge of 1 the number 1 is assumed

3) Do a “ ” with the combining capacities.

a. The 1 from Cl moves to a subscript on the Ca, and the 2+ on one Ca ion moves to the subscript on the Cl ion.

4) Write the number of needed as a subscript. Note: Cations get written first.

a. CaCl2

Transition Metals

Each of these metals can form more than one cation. Eg. Fe can form Fe2+ and Fe3+. They are located in the region of the periodic table.

Classical Naming:

1) Use the name for the atom. Fe is ferrum in Latin.

2) a. Use suffix “ ” for the lesser charge ion. Fe2+ is ferrous.

b. Use suffix “ ” for the greater charged ion. Fe3+ is ferric.

Eg. Ferr Sulfate would contain the Fe2+ ion.

Stock Naming System:

Use numerals to show the charge on the cation.

Eg. FeCl3 is iron ( ) chloride

Naming Compounds

Ionic Compounds

Molecular Compounds

Binary / Polyatomic
-Ends in “ide”
-metal is named first
-Balance the number of charges.
Ca2+ and Cl- → CaCl2 / -Ends in “ate”
-Charge is on a group of atoms.
-Balance the number of charges.
NH4+ and S2- → (NH4)2S / -Ends in “ide”
-element from left on chart is named first
-use number prefixes
-mono - 1
-di - 2
-tri - 3
-tetra - 4
-penta - 5
-hexa - 6
Eg) N2O3 – dinitrogen trioxide

Transition metal

-Find the charge on the metal, using the known charge on the non-metal.
-Classic Naming ↓ charge “ous”, ↑ charge “ic”
-Stock Naming
Eg) FeCl2 has Fe2+ - ferrous chloride or
Iron (II) chloride.
Eg) FeCl3 has Fe3+ - ferric chloride or
iron (III) chloride.

Naming Polyatomic Ionic Compounds

Poly –

Polyatomic Ions

- Ions which have one type of element in them.

- the ionic charge is on the entire of atoms.

- Ex) NO3- is a group of atoms which has an extra .

- Note: the nitrogen and oxygen atoms have bonds.

- Common polyatomic ions:

o ammonium, NH4+

o nitrate, NO3-

o sulfate, SO42-

o carbonate, CO32-

o phosphate, PO43-

o hydroxide, OH-

o perchlorate, ClO4-

Reading the name from the formula:

1) In compounds where the anion is polyatomic, the names end in “ ”.