Significant Digits

Rules:
1. all non-zero digits are significant
2. all leading and following zeros are not significant
3. all zeros between two other digits are significant
4. any zero that follows a digit and is to the right of the decimal point is significant

When performing multiplication or division, the value with the fewest number of significant digits is the number of significant digits used in the final answer

When performing addition or subtraction, the value with the fewest number of digits after the decimal place represents the number of decimal places in the final answer

Scientific Notation: used for numbers greater than 1000 or less than 0.001

To use, move the decimal to be placed after the first non-zero digit
Count the number of decimal places moved
If the decimal has moved to the left, the exponent is positive

Trends of the Periodic Table

Trend 1 Atomic Radius:

- measured in picometers (1 X 10-12 m)
- distance from the nucleus to the valence orbital
- going right and up, the atomic radius decreases
- this is because going right, there is a greater force of attraction between the protons and electrons which makes a smaller radius
- going up the periodic table, there are less valence orbitals meaning the radius becomes smaller

Key Terms:

Shielding Effect: the effect of filled inner electron orbitals on the attractiveness of valence electrons to the nucleus

Effective Nuclear Charge: the apparent nuclear charge experienced by valence electrons, result of the shielding effect

Trend 2 Ionization Energy:

- amount of energy required to remove an electron
- going right and up on the periodic table will increase the ionization energy
- this is because fewer valence shells and more protons will make it require more energy to remove electrons

Trend 3 Electron Affinity:

- energy absorbed or released when an electron is added to a neutral atom
- going right and up the periodic table, the electron affinity will increase

Trend 4 Electronegativity:

- indicator of the relative ability of an atom to attract shared electrons
- smaller atomic radius usually means greater electronegativity
- going right and up the periodic table, electronegativity increases

Types of Bonds

Ionic:

- occurs between a metal and non-metal
- two ions bond to each other
- change in electronegativity is 1.7 – 3.3

Polar Covalent:

- occurs when non-metals share electrons
- has a positive and negative end
- change in electronegativity is 0.5 – 1.7

Covalent:

- occurs when non-metals share electrons
- does not have a positive/negative end
- change in electronegativity is 0 – 0.5

Properties of Ionic and Covalent Bonds

Ionic:

- high melting point
- does not conduct electricity as a solid
- conducts electricity in a solution
- high solubility in water
- solid state
- brittle, hard consistency as a solid

Covalent

- low melting point
- does not conduct electricity as a solid or in a solution
- not soluble in water
- non-polar compounds soluble in non-polar compounds
- polar compounds soluble in polar compounds
- can be solid, liquid, or gas
- soft, flexible, waxy consistency as a solid

Lewis Structures

- symbol of element is the center
- valence electrons are represented by dots surrounding the symbol of the element
- when dealing with ionic compounds, arrows are used to represent the loss/gain of electrons
- when dealing with covalent compounds, circles around electrons represent the sharing of electrons
- line structures simply use lines to show the sharing of electrons

Nomenclature/Balancing Equations

- will not be a major part of the test, so does not need to be reviewed too much

Types of Reactions:

Synthesis

- two elements combine to create a binary compound
- universal metal + element  compound
- multivalent metal + non-metal  compound
- non-metal + non-metal  molecular compound
- non-metal oxide + H2O will create an acid – will comprise of a hydrogen cation and a polyatomic anion containing oxygen
- e.g. SO2 + H2O  H2SO3(aq) (sulphic acid)
- metal oxide + H2O  metal hydroxide (base)
- e.g. Li2O + H2O  2LiOH(aq)

Decomposition:

- compound is broken down into simpler ones or elements
- binary compound decomposes into elements
- metal nitrate decomposing into metal nitrite + O2
- metal carbonate decomposing into a metal oxide + carbon dioxide

Combustion

- chemical reaction of a substance with oxygen producing one or more oxides, heat, and light
- inorganic combustion general formula: X + O2  XO
- organic combustion only occurs with hydrocarbons which are compounds composed of only hydrogen and carbon
- complete combustion of a hydrocarbon general formula: CxHy + O2  CO2 + H2O
- incomplete combustion of a hydrocarbon general formula: CxHy + O2  C(s) + CO + CO2 + H2O

Single Displacement

- chemical reaction in which one element in a compound is displaced by another element
- always refer to reactivity chart when dealing with single displacement reactions
- metal displacing another that is below it on the activity series in an ionic solution
- e.g. Cu + 2AgNO3  Cu(NO3)2 + 2Ag
- metal ion displacing hydrogen in an acidic solution
- e.g. Mg + 2HCl(aq)  MgCl2+ H2
- active metal + water  metal hydroxide + H2
- e.g. 2Rb + 2H2O  2RbOH + H2
- halogens displacing other halogens

Double Displacement

- chemical reaction in which positive ions of two ionic compounds exchange places, creating 2 new ionic compounds
- neutralization reaction: acid + base  water + “Salt” (ionic compound)
- e.g. H2SO4(aq) + Ca(OH)2(aq)  H2O + NaCl
- precipitation reaction: reaction with two ionic solutions that creates a precipitate
- e.g. 2LiCl(aq) + Pb(NO3)2  2LiNO3 + PbCl2

Solubility

Effect of Temperature on Solubility

- every unique pure substance has its own unique solubility based on the types of bond present
- units used to describe/measure solubility is: mass of solute/100mL of solvent

Solubility of Solid
- trend: solubility of solids increase as temperature increases
- energy is required to break apart bonds of solids when dissolved in water
- as temperature increases, there is more energy to break these bonds

SATP/STP: standard ambient temperature and pressure

S = soluble
SS = slightly soluble
I = insoluble

Solubility of Gases
- trend: solubility of gases decrease as temperature increases
- with more energy, the gas particles escape the solution

Solubility of Liquids
- trend: solubility of liquids is not affected by the temperature
- solute: liquid with less amount
- solvent: liquid with greater amount

Water
- universal solvent
- small size, highly polar nature, and the capacity to form hydrogen bondsmakes water very successful at dissolving solutes
- water has a permanent dipole
- the negative end is attracted to the positive end, causing a special type of attraction called hydrogen bonding

Hydrogen Bonding
- any substance containing hydrogen and oxygen/fluorine/nitrogen
- doesn’t create an actual bond, uses strong intermolecular forces to create a force of attraction
- hydrogen bonded compounds are likely to dissolve in water

Properties
- water is held together by covalent bonds which re stronger than hydrogen bonds
- however, the hydrogen bond is stronger than regular dipole-dipole attractions
- this results in higher boiling points because more energy is required to break apart these bonds
- hydrogen bonds also result in higher surface tension

Volume/Volume Concentration (V/V) %
- volume of solute(mL)/volume of solution(mL) X 100%

Mass/Volume Concentration (m/V) %
- mass of solute(g)/mass of solution(mL) X 100%

Mass/Mass concentration (m/m) %
- mass of solute(g)/mass of solution(g) X 100%

Very Low Concentrations
- parts per million (ppm) = mass of solute/mass of solution X 106
- parts per billion (ppb) = mass of solute/mass of solution X 109

Molarity
- molar [ ]
- moles of solute/1L of solution

Dilutions
- reducing concentration of a solute by adding additional solution to the mixture
- standard/stock solution: one where the [ ] is known
- c1V1 = c2V2
- c1 is the initial [ ]
- c2 is the final [ ]
- V1 is the initial volume
- V2 is the final volume

Double Displacement Reactions
- 2 possible outcomes
- Compounds remain as ions and no reaction occurs (NR)
- New compounds created that consist of 2 of the following: solid precipitate, gas, or water

Net Ionic Equations
- an ionic compound dissociates in water and is broken up into its constituent ions
- the above occurs before a double displacement reaction happens
- net ionic equation only contains the new product and the constituents that produce this compound
- spectator ions are any ions not involved in the creation of the new product

Acids
- sour taste
- no texture
- conducts electricity in an aqueous solution
- pH less than 7
- turns litmus paper red
- phenolphthalein is colourless
- Acid + Metal  H2(g)
- Acid + Carbonate  CO2(g) + H2O(l)

Bases
- bitter taste
- slippery texture
- conducts electricity in an aqueous solution
- pH greater than 7
- turns litmus paper blue
- turns phenolphthalein pink

Arrhenius Theory
- an acid is any substance that will ionize in water to produce H ions
- a base is a substance that will dissociate in water to produce OH ions
- H ions cannot exist alone, and thus exist attached to H2O, creating hydronium: H3O
- Only valid for reactions in water

Bronsted-Lowry Theory
- acids are substances that have an H ion removed
- conjugate base is paired with the acid, and becomes the new base
- bases are substances that have an H ion added
- conjugate acid is paired with the base and becomes the new acid

Strong Acids
- will completely dissociate

Weak Acids
- will only have some of the solution dissociated
- indicated in a chemical equation by a double arrow

Monoprotic Acid
- can only give up 1 H ion

Diprotic Acid
- can give up 2 H ions (H2SO4, H2CO3)

Triprotic Acid
- can give up 3 H ions (H3PO4)

pH and pOH
- pH = -log[H or H3O]
- [H or H3O] = 10-pH
- pOH = -log[OH]
- [OH] = 10-pOH
- pH + pOH = 14

Neutralization
- Acid + Base  Salt + Water
- Titrations are done to determine the number of moles when the number of moles of H and OH are equal
- Equivalence point: the point when titration is complete (H = OH)
- End point: a sudden change occurs during a titration
- Equivalence is theoretical and determined by calculations
- End is experimental and determined by indicators

Gases

- Kinetic Molecular Theory: explains the behavior of gases in terms of random motion with no volume and no forces of attraction.
- UP temperature -> breaks bonds faster -> transforming states

Spaces Attraction Motion TypeShapeVolumeSolid Very tight Strong VibrationalFixedFixedLiquid Further Apart Somewhat Translational, rotational, vibrationalContainerFixedGasVery far apart Weak Rotational, translational, vibrationalContainerChanges

Pressures
- SI Unit = 1 Pascals (Pa) or Kilopascals (KPa)
- Older units: torr = 1 mm of Hg
- STP = 760 mm Hg = 760 torr = 1 atm = 100.3 KPa
- SATP = 250C at 100 KPa
Temperature
- Normally in Degrees Celsius, but may have to convert to Kelvin
- 1 Kelvin = -273.150C
- Convert C to K -> + 273.15
- Convert K to C -> – 273.15

Gas Laws
- Boyle’s Law
  - Pressure UP = Volume DOWN
  - P1*V1 = P2*V2
Charles’ Law
- Temperature UP = Volume UP
- V1/T1 = V2/T2
Gay-Lussac’s Law;
- Temperature UP = Pressure UP
- P1/T1 = P2/T2
- Combined Gas Law
  - Accounts for pressure, temperature, and volume changes
  - (P1*V1)/T1 = (P2*V2)/T2
  - Ideal Gas Law properties
Molar volume of ideal gas at:
- STP: 22.6L/Mol
- SATP(SAP): 24.8L/Mol
- Avogadro’s Law
  - States that ALL ideal gases at the same temperature and pressure contains the same number of molecules of gases
  - n1/V1=n2/V2
  - Ideal Gas Law
    - PV = nRT
      - Where R is the universal constant of gases 8.314
      - Temperature always in K
      - Volumes in L
      - Pressure in KPa usually recommended
Deriving Ideal gas law
- Molar Mass
  - mRT/PV = M (g/mol)
  - Density
    - m/v = MP/RT (g/L)
  - Molar Concentration
    - C = n/V = P/RT (mol/L)
    - Pressure of Water Vapors
      - Pdry gas= Patm – Pwater vapour
      - Pressure of water vapour obtained with a table that relates it to temperature
      - Gas Stoichiometry

Balancing Equation

Convert everything to moles (from formulas etc.)

Mole to mole ratio

Solve for your unknown

Must find limiting reactant if necessary

Volume to Volume: nothing to do since volume = number of moles

STOICHIOMETRY

Isotopes and average atomic mass
- Isotope: element variations with different atomic mass but same atomic number
- Isotopic Abundance: the relative amount in which each isotope of an element
- Calculating: given 2 isotopes of an element [B: 10.01u] [B: 11.01u]
  - 1) Given average atomic mass: 10.81u
    - If given percentages, you can find average atomic mass by adding each portion
- 2) Set variables for unknown values
  - Let x represent % abundance for 10/5 B
  - Let 1-x represent % abundance for 11/5 B
- Calculate:
  - 10.01x+(1-x)11.01=10.81
  - 10.01x+11.01-11.01x=10.81
  - -1.00x=0.2
  - X=0.2
  - The Mole and Avogadro’s constant
    - Mole: the SI unit that is used to measure how much of a substance. (balanced eqn coefficient are moles)
    - Converting Moles -> number of particles
      - Number of Moles = number of particles / Avogadro’s constant
      - Avogadro’s constant: 6.02214179 x 10^23
Converting Moles to Mass/molar mass
- Molar Mass: how much grams per mol of an element = atomic mass
- Number of Moles = mass / Molar Mass
- Percentage Composition
  - Laws of definite proportions: a law stating that a chemical compound always contains the same proportion of elements by mass
  - Mass percentages (H in H2O) = molar mass of H (1.01) / molar mass of H2O (18.02) = 11.2%
  - Empirical Formula
    - A formula that shows the smallest whole number ratio of elements within a compound
    - Lactic Acid: a chemical name
      - Empirical Formula: CH2O
      - Molecular Formula: C3H6O3
Calculations: when given percentage composition of different elements
- 1) Find the number of moles for each element (assume % is per 100g for simplicity)
- 2) Divide all the moles by the smallest number mole to reduce the decimals
- 3) If not whole numbers yet, multiply by some special factors to reduce decimals
  - 1.25 x ¼ = 1
- 4) Rewrite equation into the empirical formula
- Molecular Formula
  - A chemical formula that expresses the exact ratio of atoms (whole numbers) for a compound
  - NO REDUCTION REQUIRED, and not suitable for ionic bonds (metal + non metal)
  - Determining Molecular Formula:
    - Molar mass of compound must be given
    - Let X represent a multiplier that relates the empirical formula to molecular formula
    - X = molar mass of actual compound / molar mass of EF
    - Multiply all the elements in the empirical formula by that X multiplier
    - Stoichiometry: study of relationships between reactants and products involved in chemical reactions
      - IOWs, study of how much of reactants are required to produce sufficient products
      - Simply use number of moles in reactions (EF) to become their ratios between each other
        Al + 3 CuSO4 -> Al2(SO4)3 + 3 Cu : 2:3:1:3 ratio
To determine amounts, use ratio and plug in numbers correctly
Limiting/excess reactants
- Limiting:Reactants that are not completely used up in a reaction
  - When this reactant is used up, the reaction stops
Excess Reactants:the reactant that is left over after a reaction is over
- Regardless of how much excess remains, the reaction will stop once the limiting is used
When 2 or more reactants seem limiting, one must pick the reactant that is limiting all the others
Mass to mass Stoichiometry
- Converting mass through stoichiometry then to another mass
- Process:
  - 1) Balance equation
  - 2) Convert mass to moles
  - 3) perform mole to mole ratio
  - 4) convert new mole back to mass
  - Finding Hydrate
    - Given a hydrate, and when hydrate is evaporated, and new mass is taken, how much of the hydrate was there?
    - Process
      - 1) find mole of non-hydrate portion
      - 2) find mass of the water that evaporated
      - 3) find mole of H2O
      - 4) based on ratio, find out how many mols of water should exist
        
        Based on the amount of amount of actual results in ratio with theoretical ratio
        
        Find the amount that were theoretical based on regular ratios of stoichiometry
        
        (OR: divide by the smallest mole to get rid of as much decimals)
        
        Percentage Yield
        How much percent of the actual product will happen over the theoretical amount
        Percentage Yield = Experimental yield/theoretical yield x 100%
Factors:
- Competing Reaction
  - Other reactions will happen at the same time, limiting the reactant produced
- Reaction Rate
  - Slow reactions take longer than time allocated to complete, collected immaturely
- Purity of reactant
  - Contaminants in the reaction are unaccounted for in the initial calculations
- Laboratory Techniques
  - Improper techniques leads to loss of reactants and reactants
  - Soluble solvent in filtrate instead of filter paper
  - Soluble products washed away in rinsing
  - Instruments used weren’t accounted for
When looking for masses
- Look for ratio of compound to water
- Compare ratio with amount of solvents used, there’s the grams
- Nuclear reactions
  - Involves changes to the nucleus of the reactants
    - Changes the number of protons or neutrons
3 Types of nuclear reactions
- Alpha ParticleDecay (a decay)
  - Loss of 2 protons and 2 neutrons
  - Always creates 4/2 He product
- Beta Particle Decay(b decay)
  - Loss of an electron-like beta particle
  - 0/1 e always a product, adds one neutron to other product
- Gamma Decay(Y decay)
  - Releases beta particle as well as gamma energy
Conditions of nuclear reactions
- Sum of mass numbers of product = sum of mass numbers of reactants
- Sum of atomic numbers of product = sum of atomic numbers reactants
Nuclear Fission
- Occurs when an unstable isotope splits into smaller fragments. Multiple neutrons released along with high energy (radiation)
Nuclear Fusion
- Nucleus absorbs particles, merging 2 atomic nuclei together

	Created by luxshiga_1997 over 10 years ago

Note on SCH3U Chem Exam Review, created by luxshiga_1997 on 01/15/2014.

SCH3U Chem Exam Review