Chem final

Ion

an atom or molecule with a net electric charge due to the loss or gain of one or more electrons.

Cations

Positively charged ions (lost electrons). Usually metals.
Example: Na⁺, Ca²⁺

Anions

Negatively charged ions (gained electrons). Usually nonmetals.
Example: Cl⁻, O²⁻

Transition Metals with More Than One Charge

• Use Roman numerals to indicate the charge.

  • Iron(II) = Fe²⁺, Iron(III) = Fe³⁺

  • Copper(I) = Cu⁺, Copper(II) = Cu²⁺

Polyatomic Ions (great 8)

Highlighted on PT

Ionic Compounds

substances composed of oppositely charged ions (cations and anions) held together by electrostatic forces, which are known as ionic bonds

Naming Binary Ionic Compounds (metal + nonmetal)

• Name the metal (cation) first.
  Name the nonmetal with the  -ide ending.

  • Example: NaCl = sodium chloride

Naming with Transition Metals

• Use Roman numerals for the metal’s charge.

  • Example: FeCl₃ = iron(III) chloride

Writing Formulas

• Balance the charges to make a neutral compound.

  • Aluminum oxide: Al³⁺ + O²⁻ → Al₂O₃

Polyatomic Compounds

• Keep the polyatomic ion name unchanged.

  • Ca(NO₃)₂ = calcium nitrate

  • Mg(OH)₂ = magnesium hydroxide

prefixes

1. mono- (1) (not used on first element)

2. di- (2)

3. tri (3)

4. tetra (4)

5. penta- (5)

6. hexa- (6)

7. hepta- (7)

8. octa- (8)

9. nona- (9)

10. deca(10)

when do you use the prefixes?

When naming molecular compounds, especially those composed of two or more nonmetals

• Examples

  • CO = carbon monoxide

  • CO₂ = carbon dioxide

  •  N₂O₅ = dinitrogen pentoxide

Acids

• Starts with H in the formula

  • Naming Acids

• If it contains -ide anion, it is hydriodic acid

  • HCl = hydrochloric acid

• If it contains -ate → __ic acid

  • HNO₃ (nitrate) = nitric acid

• If it contains -ite, __ous acid

  • HNO₂ (nitrite) = nitrous acid

Know how to write a chemical equation from words

• Identify the chemical formulas for all compounds mentioned.

• Write the reactants on the left and the products on the right.

• Use symbols:

  • (s) solid, (l) liquid, (g) gas, (aq) aqueous

  • Example:
    "Hydrogen gas reacts with oxygen gas to form water."
    → H₂(g) + O₂(g) → H₂O(l)  (Balance it later)

Be able to balance equations with coefficients

1. Count the atoms of each element on both sides.

2. Add coefficients, whole numbers placed in front of reactants or product

3. Do NOT change subscripts!

• Example:
  Unbalanced: H₂ + O₂ → H₂O
  Balanced: 2H₂ + O₂ → 2H₂O

How many types of general reactions are there?

5

Synthesis (Combination)

•  A + B → AB

• Ex: 2Na + Cl₂ → 2NaCI

•  Combine elements or compounds.

Decomposition

• AB → A + B 

• Ex: 2H₂O → 2H₂ + O₂

• Split a compound into simpler substances.

Single Replacement

• A + BC → AC + B

• Ex: Zn + HCl → ZnCl₂ + H₂

•  Use the activity series to see if replacement happens.

Double Replacement

• AB + CD → AD + CB

• Ex: Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCI

• Swap ions and check the solubility rules

Combustion

 Hydrocarbon + O₂ → CO₂ + H₂O

• Ex: CH₄ + 2O₂ → CO₂ + 2H₂O

• produces CO₂ + H₂O if complete.

Know 3 ways in which matter is measured

• By Mass (grams)

  • Count by tens to an individual number of items 

  • Measured using a balance.

  • Used in labs to weigh substances.

• By Moles

  • Moles by weighing

  • Based on the number of particles (atoms/molecules).

  • 1 mole = 6.022 × 10²³ particles.

• By Volume

  • Volume by finding the space it takes up 

  • For gases at STP (standard temperature and pressure).

  • 1 mole = 22.4 L at STP

Know Avogadro’s number and be able to define a mole

• 6.02 × 10²³ particles/mole

• A mole is a counting unit like a “dozen” but equals 6.02 × 10 23 particles.

Be able to calculate molar mass and convert between moles and mass (grams)

• Molar Mass - The mass of 1 mole of a substance in grams/mole (g/mol).

• Equal to the sum of atomic masses from the periodic table.

• Example:
  H₂O:
  H = 1.01 × 2 = 2.02
  O = 16.00
  → Molar mass = 18.02 g/mol

Convert between moles and particles, or volume

• To convert between moles and particles

  • Use 6.02x10^23 particles = 1.0 mol conversion factor

  • Ex. 3.2 mol =? Particles

    • 3.2 mol x 6.02x10^23/1 mol = 1.9x10^24 particles

Find % composition

• Add up the masses to the total mass

• Divide each element's mass by the  total

• Multiply by 100%

• If a sample is found to contain 5.4% oxygen and 8.2% sulfur, what is the % composition

  • 5.4+8.2 =13.6

  • 5.4/13.6 = 39.7

  • 8.2/13.6 = 60.3

    • If you know the formula, you can find the percent composition by mass

• Percent comp = mass of element of a compound x 100

• Find the molar mass of the total and divide each element by the total, x by 100

  • Molar mass: 18

  • 2/18x100 = 11.19

  • 16/18x100 = 88.889

Calculate both empirical and molecular formulas

• Empirical

• change % sign to grams

• divide by mass of element to find moles

•  divide by smallest mole amount to find ratio

• Molecular: Divide given molar mass by emp molar mass to find whole number to multiply emp form by

Know how to tell quantities of reactants or products like a recipe

• Coefficients tell you how many moles of each substance react or are produced.

• Example: 2H₂ + O₂ → 2H₂O 2H₂ + O₂ → 2H₂
  This means:

  • 2 moles of hydrogen react with 1 mole of oxygen to make 2 moles of water

Convert between moles and moles

• Use coefficients as a conversion factor

Convert between grams and grams

• Convert to moles first, then grams g given to mol given to mol want to g want

Know how to find the % yield.Know how to find the % yield.

= actual/theoretical × 100

Know the 3 parts to the kinetic theory.

• Particles are small, have insignificant volume, are far apart, and are independent.

• The motion of gas is rapid, chaotic, random, and constant for a given temperature p 

• Collisions with another particle or object are perfectly elastic (no energy lost)

Gas Pressure

when gas particle collides w/ something

Atmospheric Pressure

• collisions of air molecules w/ object

evaporation

Liquid → Gas (slow, happens at surface)

Condensation

Gas → Liquid

Boiling

• Liquid → Gas (fast, throughout the liquid at the boiling point)

Condensation

(Gas → Liquid)

Sublimation

Solid → Gas (skips liquid phase, e.g., dry ice)

Deposition

Gas → Solid, e.g., frost

Melting

Solid → Liquid (e.g., ice melting)

Freezing

Liquid → Solid

Know why some substances are solids and some are liquids, and spare me gases at the same temperature

• Depends on Intermolecular attractions; the stronger IM attractions, the lower the vapor pressure

Be able to explain hydrogen bonding

• Hydrogen Bonding-  when H bonds w/ unshared pair of electrons on O from a different molecule, causes unique properties of water

• surface tension is the inward pull that minimizes surface area

• can be broken with surfactant

• Vapor Pressure is low and slow evaporation

• Heat capacity is high, heats and cools slowly

Solid Water is less dense than liquid(4⁰C most dense)

Tell why water is unique in surface tension, vapor pressure, boiling point, and density.

• Surface Tension- tends to hold drops in a round shape, thin skin on surface

• Vapor Pressure is lower with solute for simple liquids and slowly evaporates

• Boiling Point rises when a solute is present

• Liquid Density is higher than solid water

Solute

The substance being dissolved (e.g., salt).

Solvent

• The substance doing the dissolving (e.g., water).

Polarity

• "Like dissolves like”

• polar dissolves polar

• Nonpolar dissolves nonpolar

Water

univerisal solvent, polar

Dissociation

• when the ionic compounds split into smaller particles

Dissolving

• when a solid or liquid solution is formed in a solvent

Know the 3 factors to get something to dissolve faster

• Increase temperature → particles move faster.
  Stir or agitate → more interaction between solute and solvent.
  Particle size (crush the solute) increases surface area.

Be able to work on solubility problems.

Grams per 100g of water, use chart

Calculate the Molarity of a solution

• Formula: M = moles of solute/liters of solution

• Molarity is the amount of a substance in a certain volume of solution

Compare concentrated and dilute solutions

• Concentrated: High amounts of solute 

• Dilute: small amounts of solute

• You can dilute a solution by adding more solvent

vapor pressure depression

• VP lowers when solute is present

freezing point depression

• FP lowers when solute present

• -1.86°C per mol in 1L of water

boiling point elevation

• BP raises when solute present

• .51⁰C per mol in 1L of water

Properties of acids

• Tastes sour or tart (like lemon juice)

• Are electrolytes (conduct electricity when in water)

• Causes litmus to turn red or pink 

• React with many metals to release H₂.

Properties of bases (also called Alkaline)

• Tastes bitter (ex-soap)

• Feels slippery 

• Are electrolytes

• Turns litmus paper blue

• Not found in food, but found in cleaning supplies

Arrhenius Acid

• contains Hydrogen that become a H+ ion in water, must be joined to more electronegative atom

Arrhenius Base

•  contain hydroxide that can become ion in water (OH-)

Bronsted Acid

• hydrogen ion, Donor

Bronsted Base

• hydrogen ion, Acceptor 

Be able to do calculations with K_w

• The ionization constant of water (Kw) represents the equilibrium constant for the autoionization of water. It is defined as:

•  [H+] x [OH-] = 1x10^-14M

Know the pH formula and scale, and be able to do calculations

• The pH is a measure of the acidity or basicity of a solution

• pH = - log[H+}
  pH is a scale that ranges from 0-14

  • 0-7= acidic (closer to 0= more acidic)

  • 7-14= basic (closer to 14 = more basic)

  • 7 - neutral (not acidic)

Know how to complete a Titration problem

• Used to find the unknown concentration of an acid  or base by using a certain amount and concentration

• Regardless of concentration or amount, the number of moles must be equal

• An indicator (color change) is used to determine the endpoint (neutralization) of the titration

• Ex 25.0 mL of an acid (HXL) with an unknown concentration is titrated with 16.8 mL of a 1.25M solution of NaOH. What is the concentration of the acid?

  • Since it is ml, we need to move the decimal point.

  • 1.25 m=x/0.0168

  • x = 0.021 mol acid

  • Acid = 0.021 mol/0.025 = 0.84M

  • Ma(va) = Mb(Vb) 

Know how to complete a neutralization reaction

• Acids and bases cancel each other out when added together (neutralize)

• Products are water and salt (not necessarily table salt)

• Need to have the right amounts of each

  • A small weak base won't completely neutralize a large amount of strong acid.

  • Ex. HCL + KOH → HOH + KCL (DR 1:1:1:1) (both are strong)

  • Ex. 2HNO₃ + Ca(OH)₂ → 2HOH + Ca(NO₃) ₂

    • For the 2HOH, we don't have 2 h with OH because it is with the O