Chapter 4 Chemistry Exam

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1

pH of an Acid

pH less than 7

2

pH of a Base

pH more than 7

3

Litmus color of Acid; taste of Acid

Red litmus; sour taste

4

Listmus color of Base; taste of Base

Blue litmus; bitter taste

5

Acid's Dissociation in Water

H+ Dissociation

6

Base's Dissociation in Water

OH- Dissociation

7

Strong Acids

Hydroiodic Acid (HI), Hydrobromic acid (HBr), Hydrochloric acid (HCl), Perchloric acid (HClO₄), Sulfuric acid (H₂SO₄), Nitric acid (HNO₃)

8

Strong Bases

Sodium hydroxide (NaOH), Lithium hydroxide (LiOH), Potassium hydroxide (KOH)

9

Metal Oxides in Water Produce...

Bases
ex. CaO + H₂O → Ca(OH)₂

10

Neutralization Reaction

Applies to all strong acid and base reactions because the solution is neither acidic or basic in the end.
Water or salt can be formed.

11

Commonly formed gases in Gas-forming Reactions

CO₂, H₂S, SO₂, NH₃

12

Difference between Solvent, Solute, and Solution

Solution=Solute+Solvent
example of this: saltwater (solution)=salt (solute)+ water (solvent)

13

Traits of Non-electrolytes

Dissolve without breaking into pieces.
No charged particles; solution doesn't conduct electricity.
Are molecular compounds.
examples: Sugar (glucose) C₆H₁₂O₆, Ethanol CH₃CH₂OH, Ethylene glycol C₂H₆O₂

14

Traits of Electrolytes

Splits into many charged particles when dissolved.
Ionic compounds that are very soluble mostly split apart into ions.

15

K.I.S.S. Rules: Soluble Cations

Sodium ion Na⁺, Ammonium NH₄⁺, Potassium ion K⁺

16

K.I.S.S. Rules: Soluble Anions

Nitrate NO₃⁻, Acetate CH₃COO⁻, Chlorate ClO₃⁻, Perchlorate ClO₄⁻

17

Silver Group

Silver Ag, mercury Hg, and lead Pb

18

K.I.S.S. Rules: Mostly Soluble

Chloride Cl⁻, Bromide Br⁻, Iodide I⁻, except with Silver Group: silver Ag, mercury Hg, and lead Pb

19

K.I.S.S. Rules: Sometimes Soluble

Sulfate SO₄²⁻, except with Silver Group: silver Ag, mercury Hg, and lead Pb, and with Barium Ba, and Strontium Sr

20

K.I.S.S. Rules: Everything Else

INSOLUBLE.

21

Redox (Oxidation-Reduction) Reaction

A reaction in which electrons are transferred from one substance to another.

22

Oxidizer

Attracts and takes electrons from another substance, it's charge goes down; it is reduced.
OXygen is a good OXidizer (hint hint), and so are the elements around it on the periodic table.

23

Reducer

Gives up electrons when they are attracted by another substance, it's charge goes up; it is oxidized.
Li is a good reducer, and so are the elements around it on the periodic table.

24

Oxidation Number

Rates how well an atom takes or gives electrons in comparison to others.

25

Oxidation Number of a Pure Element

Zero

26

Oxidation Number of Monatomic Ions

Equal to the charge on the ion

27

Oxidation Number of Fluorine (Fl)

-1 in compounds; it's the best at stealing electrons.

28

Oxidation Number of Chloride Cl⁻, Bromide Br⁻, Iodide I⁻

-1, except against Oxygen and Fluorine

29

Oxidation Number of Oxygen and Hydrogen

H has a 1+, O has a 2-

30

Rule for Oxidation Number of Compounds

The sum of oxidation number in a compound must equal the overall charge of the compound.

31

Redox Reaction occurs when...

The Oxidation Number of an element changes during the reaction; it is reduced or or oxidized.

32

Weak Electrolytes

Only ionize to a small extent. Acetic Acid CH₃CO₂H and Ammonia NH₃ are examples.

33

Strong Electrolytes

Strong acids and strong bases are strong electrolytes.

34

Steps to Solving Net Ionic Equations

1. Write the equation.
2. Label charges and states of matter.
3. Determine spectator ions and the precipitate
4. Eliminate spectators from both sides
5. Balance the final equation

35

Diprotic Acid

An acid containing two ionizable hydrogen atoms per molecule. Sulfuric Acid H₂SO₄ is an example.