Chem 11 Jan

Chemistry 11 January Exam

Unit 1 - The Processes of Chemistry

1.2 Significant Figures

Measuring

• Science involves measuring

• No measurement is perfect (exact) because we are limited by:

  • Imperfections in the measuring device

  • Imperfections of the measurer (you)

Reading a scale

• When reading a scale, record digits until you get to one that you are not quite sure of (between the smallest lines)

• Record it and no more

• This gives you a number with 1 uncertain digit

• For electronic balances, record what shows up and nothing more or less

1.3 Uncertainty

To find the average of several tests:

1. Find the average (mean)

2. Round it to the uncertain digit

   1. The uncertain digit should have 3 or more different digits to be considered as one

3. Find the range (number furthest away from the average)

4. Format as: average ± range/uncertainty unit

   1. E.g. 1.37 ± 0.05 V

• Uncertainty is a 1 digit number that indicates the range in which the accepted or true value most likely lies

• Uncertainty represents the natural variation that is inherent in measurement

• All measured values (those with units), have uncertainty

1.4 Math Tools in the Lab

Accuracy

• How close a measurement is to the accepted value (CORRECT)

Precision

• How close a series of measurements are to each other (CONSISTENT)

• “Precisely wrong” means getting consistent results, but getting them wrong

• Good accuracy and poor precision is not possible at all

Percent Error

• Indicates accuracy of a measurement (ie how far your measurement is away from the accepted value)

*for percent error, look at the result of the numerator with the absolute values included, and that should be the number of significant figures your percent error should have (typically 1 SF)

Scientific Notation

• Move decimal until there is 1 digit to the left, places moved = exponent

  • Large number (>1) → positive exponent

  • Small number (<1) → negative exponent

• Only include sig figs (all digits before the multiplier are significant

*notes: use shortcut button on calculator (x10x button)

Proportions

• Direct proportion

  • y ∝ x

  • As x increased, y increased

• Inverse proportion

  • y ∝ 1x

  • As x increased, y decreased

  • Greater the denominator, smaller the function

Significant Figures

• Indicate precision of measurements only

• There are 3 kinds of numbers

  • Measured numbers

    • E.g. 3.24m, 180kg, etc

    • Need SFs

  • Counted numbers

    • E.g. 4 coins, 16 girls, etc

    • Infinite SFs, no uncertainty

  • Definition or conversion

    • E.g. 1 m = 100 cm

    • Infinite SFs, no uncertainty

• Recording SFs

  • SFs in a measurement include the known digits and a final estimated digit

• Counting SFs

  • Count all numbers except for

    • Leading 0’s

      • E.g. 0.0025

    • Trailing 0’s without a decimal point

      • E.g. 2500

  • These zeroes are considered place holders and are not significant

  • Provide zero information about the precision of the measurement (only information it  gives is how big or small the measurement is)

• The number of meaningful digits in a measurement includes 2 uncertain digit

• Sometimes in order to express the amount of SFs it needs to be written in scientific notation

1.5 Calculations with Significant Figures

Addition & Subtraction

• The answer should have only as many figures as the least precise number in the problem

• The final measurement is limited by the less precise number

• With scientific method you can either:

1. Expand both & add/subtract (harder with large exponents)

2. Rewrite one of the numbers’ scientific notation so both have the same exponents

Multiplication & Division

• The number with the fewest SFs determines the number of SFs in the answer

• Ignore exact (counted & conversion) numbers (do not limit the # of SFs in the answer)

1.6 Factor Label Method

Metric system 

•  Basic units and is the standard form of units

• Base/parent unit - e.g. m, s, kg

How to convert between metric units

1. Write starting quantity with the unit

2. Multiply by a conversion factor

• Make the denominator the same unit in the numerator (or base unit & continue to convert)

• Put a “1” before the other unit

• Write the correct multiplier before base unit

1.7 Non-metric Unit Conversions

• Can use the factor label method to convert between any 2 units of which we know the conversion rate

• Ex. 1.00 kg = 2.21 lb

  • Is a conversion rate

    • Used directly in the conversion factor

• Some of these metric ⇌ non-metric conversions are measured, therefore significant figures haave to be taken account (including input/original number)

• Think definition vs. measured

Other formulas

d = mv   density unit - g/mL or g/cm³ (1mL = 1 cm³)

Rates

• One unit/another unit

• Can be flipped around to match other/desired unit

• E.g.

Unit 2 - The Nature of Matter

2.1 Matter & its Changes

Matter

• Is the stuff around us

• Is composed of particles (atoms, elements, compounds, or a mixture of these)

Physical Change

• Nothing new is produced

• Change of state (s), (l), (g), (aq)

• Dissolving in water does not result in a new substance being produced

• E.g.

(do not put water in the equation because it is not a reaction with water, refer to more details in unit 4)

Chemical Change

• Something new is produced

• New formulas (atoms are rearranged)

• New properties

• Colour changes

• Bubbles

• Precipitate (solids form in solution)

• E.g.

(E is for energy)

Physical Property

• Can be found without creating a new substance

• E.g. solution colour, density

Chemical Property

• Describes the ability of a substance to undergo chemical reactions and change into new substances, either by itself or with other substances

• E.g. flammability, reactivity

Extensive (or extrinsic)  property

• Depends on the amount of matter present 

• E.g. 

• Boiling time

• Volume

• Mass 

Intensive (or intrinsic) property

• Depends on the identity of the substance, not the amount

• E.g.

• Boiling point

• Density

• Conductivity

The Kinetic Molecular Theory (KMT)

•  Solids

  • Particles vibrate

  • Take rigid/defined shape

  • Measurable volume

  • Kinetic energy lowest, potential energy highest

• Liquids 

  • Particles slide between each other 

  • Takes the shape of the container

  • Non-measurable volume

• Gasses

  • Particles move freely with high energy

  • Take the shape of the container

  • Non-measurable volume

  • Kinetic energy highest, potential energy lowest

Changes in Physical State

Heating Curve

1,3,5: as more heat is added to the system, particles move faster, and hence possess more kinetic energy, and temperature increases

2,4: no temperature increase, heat added to the system is converted to potential energy to pull particles apart

Note: there is no temperature change during phase change

2.2 A Classifying Matter

All matter is made of particles

• Either an atom (1 type of particle), or a group of atoms bonded together

• Atoms

  • The smallest possible unit of an element which retains the fundamental properties of an element

  • Ex. silver (Ag), copper (Cu), hydrogen (H)

• Molecule

  • A cluster of two or more atoms held together strongly by electrical forces

  • Term used for covalent compounds or some elements (diatomic or polyatomic)

  • Ex. water (H₂O)

• Ion

  • An atom molecule which possesses an electrical charge

  • Ex. sodium ion (Na⁺), chloride ion (Cl^-)

Pure Substances

• A substance with a constant, unchangeable composition (contains only one type of particle)

• Ex. sugar, water, copper

• Elements

  • A substance that is composed of only one kind of atom and which cannot be separated into simpler substances as a result of any chemical process

  • Ex. can be an atom like Cu or a diatomic/polyatomic like N₂ or S₈

• Compounds

  • Any substance that is composed of more than one kind of atom

  • Only molecules or ionic compounds can be compounds like H₂O or NaCl (s)

Mixtures

• Made up of two or more pure substances, such that the relative amounts of each can be varied

• Contains more than one type of particle

• Ex. salt dissolved in water, gravel, pencil

• Heterogenous

  • A substance consisting of distinct regions separated by visible boundaries

  • Ex. a human being, gravel, pencil, oil & water

• Homogeneous 

  • A substance without distinct regions or boundaries

  • Ex. dissolved salt in water, water & food colouring, sugar solution

• 3 types:

  • Suspension (heterogeneous)

    • Larger particle size

    • Sediment upon standing

    • E.g. flour in water

  • Colloid (homogenous)

    • Has a uniform composition

    • Translucent or “milky”

    • Medium particle size

    • E.g. diluted milk

  • Solution

    • Solution = solute (minor component) + solvent (major component)

    • If the solvent is water, it is an aqueous solution (aq)

    • E.g. vegetable oil

* not heavy enough to let gravity pull it down

2.2 B Separation of Mixtures

• Density separation

  • Method 1: based on different densities

    • Mixture of A+B → add liquid → different densities layered

    • E.g.

• Method 2: centrifugation

  • Similar to density separation but with extra spinning

  • Machine that spins at high speeds to separate

• Decanting/pouring

  • Carefully pouring of the liquid and leaving the sediment in the bottom of the original container

  • Quick & easy

  • Small amount of product might be lost

• Filtration

  • To separate an insoluble solid from the mixture

  • Keeps as much product as possible

  • More complicated set-up than decanting

2 miscible liquids (homogenous mixture) 

• Distillation 

  • Separate mixtures of miscible liquids by using their different boiling points

  • Think of the wine set up

• Chromatography

  • Separate substances in a mixture by having a flowing fluid carry them at different rates through a stationary phase

  • Some compounds in the mixture adheres to the stationary phase more, hence it travels slower

  • Gas, column, thin layer, & paper chromatography

  • Retention factor

    • Flow speed relative to that of mobile phase

    • Rf=distance the substance flowsdistance the solvent flows

    • There is no unit, because the units in the distances “cancel out”

    • Higher the Rf, the less attracted to the stationary phase, collected first

2.3 Ionic Compounds

Ionic bond

• Metal + non-metal

• Very strong bond

Writing formulas

1. Write symbols with combining capacities (charges found on periodic  table)

2. Balance charges (criss-cross)

3. Write formula (write roman numeral if metal is multivalent)

E.g.

Naming ionic compounds

1. Name the metal (check its combining capacity ie roman numerals)

2. Name non-metal ending to -ide, or find polyatomic name

3. Use a roman numeral for multivalent

E.g.

2.4 Covalent Compounds

Covalent bond

• Non-metal + non-metal

• Covalent molecules can be solid (e.g. S₈), liquid (e.g. H₂O ), or gasses (O₂)

• Form molecules

• Contain covalent bonds that result from electron sharing

Naming covalent compounds

Mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

1. State prefix before first element’s name, mono is never used for the first element

2. Second element changes to -ide

3. Vowels (a, o) on prefix are omitted if followed by vowels (a, o) (e.g. carbon monooxide, carbon monoxide)

E.g. 

Hydrates

• A solid compound containing water molecules combined in a definite ratio as an integral part of  the crystal

• Ie the water is bonded with the ionic compound

• For every x unit of ionic compound, there is x unit of water molecules

1. Name the ionic compound normally

2. Use  prefixes for the # of water molecules

3. Name water molecule “hydrate”

E.g. 

Unit 3 - The Mole

3.1 Intro to the Mole

Avogadro’s number

• How heavy is 1 molecule?

• How can we go about determining the weight of a sample of molecules?

• Atoms & molecules are extremely small & light

• Can weigh large numbers of them, & divide it to find the weight of 1

Avogadro decided to take 1.00 g of the smallest atom (H) & determined how many H atoms there are in 1.00g of H

• 1.00g H = 6.021023 particles = 1.00 mole

• Note: the “6.02” part is a measured number, and has 3 SFs therefore take into account SFs in mole calculations

Note: really just a fancy name for a big number of particles

What  is considered as a particle?

• Atom (element) e.g. Cu

• Molecule (covalent) e.g. CH₄

  • Or diatomic/polyatomic

    • Most common ones are: N₂, H₂, O₂, F₂, I₂, Cl₂, Br₂, S₈

• Formula units/FUs (ionic) e.g. NaCl

Conversion Factors

1. 1.00 mole6.021023 particles (of any pure substance)

2. 6.021023 particles (of any pure substance)1.00 mole

3.2 Molar Mass

• A mole is a quantity equal to the number of atoms in the atomic mass of any element expressed in grams (ex. The number of atoms in 1.0 g H, 16.0 g O, 63.5 g Cu, etc)

• Atomic mass of an element is the relative mass compared to Hydrogen

• Unit is the atomic mass unit (amu)

• Molar mass is the mass of 1 mole of particles, unit is g/mol

• The molar mass is a new conversion factor that can convert between grams & moles

• Avogadro’s number is another conversion factor that can convert between moles & particles

• You can use this to convert between each, but more will be added later on

3.3 Mole Calculations

• Calculating between different units using conversion factors

3.4 Mole Calculations II

• Calculate x molecules or x atoms within x amount, and vice versa

• Create your own ratios etc depending on the question

• SFs not needed for these ratios, as these are counted numbers (the other ratios still have SFs, cannot ignore them entirely)

• Note: remember dimensional analysis, the desired unit on the top, the cancellable unit on the bottom

3.5 Empirical Formula

Empirical Formula

• Simplest whole number ratio between atoms in a compound, determined experimentally by measuring the mass of the elements that combine to form a compound

Molecular Formula

• Formula of the molecular unit, a multiple of the empirical formula

To find the empirical formula

1. Turn all from grams to moles, if a percentage, assume the compound is made up of 100g

2. Divide by the smallest # of moles

3. May have to multiply by a multiple if there is a decimal

3.6 Percent Composition

To find percentage composition

1. Multiply x number of an atom to its molar mass

2. Add up all of the calculated molar masses

3. Divide each element’s molar mass by the total molar mass then multiply by 100

4. Each number should add up to approximately 100, round to SFs as instructed by the question

3.7 Weighing Gasses

Avogadro’s hypothesis states that equal volumes of gasses have the same number of particles

E.g. 1.0 L of CO₂ 1.0 L of H₂

• Same number of particles (same number of CO₂ as H₂)

Ideal Gas Law

• Assumptions

1. The volume of the gas molecules can be neglected in a container

2. There is no interaction between gas molecules

Indirect/Combustion Analysis (hardest possible test question, supposedly)

Unit 4 - Chemical Reactions

4.1 - Intro to Chemical Reactions

Signs of a Chemical Reaction

• Evolution of heat & light

• Formation of a gas

• Formation of a precipitate

• Colour change

Law of conservation of mass

• Mass  is neither created nor destroyed in any chemical reaction

• Total mass stays the same

• Atoms only rearrange

Equation Symbols

Quick Review

• Atoms: singular elements

• Molecules: covalent compounds

• FUs/Formula Units: ionic compounds

4.2 Writing Chemical Reactions

Why balance equations?

• Mass is neither created nor destroyed (law of conservation of mass) in a chemical reaction

• Thus there must be the same of each thing on each side

• Note: solution means it is aqueous

4.3 - Synthesis & Decomposition

• Phase symbols (s, l, g, aq) are used to indicate phases of reactants & products

• Polyatomic elements:

• Diatomic elements:

• (both should be written on p-table)

Synthesis

• make/combine elements to form compounds

• A + B → AB

• Ex. 

• Ex. 

Decomposition

• Breakdown, compounds break down into elements

• AB → A +B

• Ex. 

• Ex. 

4.4 - Single Replacement Reactions

Single Replacement Reactions

• Occur when a single uncombined element replaces another element in a compound

• Can only swap with its own kind (metal/non-metal)

AX +BY → AY + B (metal element replaces metal ion)

X + AY → AX + Y (non-metal element replaces non-metal ion)

• Look at the activity series (separate ones for metal & non-metal)

• If the single one is above, they can replace

• If the single one is below, no reaction

Activity Series

• Higher = more reactive/active

• Lower =  less reactive/active

• Why? 

  • The formation of ionic compounds is when metals donate electron(s) to non-metal elements

  • More active metals donates electrons more readily, better at doing the job of “electron donor”

Note: for multivalent metals such as below, just choose any one of the charges

4.5 - Double Replacement Reactions

Double Replacement Reactions

• Occur when the metal ions of two aqueous compounds switch places

• Precipitates (solids) often form as a result

M1B +M2Y → M1Y + M2B

• The aqueous solutions will dissociate and some may form bonds

• Look at solubility chart, “low solubility” means that a precipitate or solid will form

• If both end up being aqueous, then there is no reaction (NR)

Acid Base Neutralisation Reactions

• A type of double replacement

• Acid + base → salt + water

• Ex. 

Acid Carbonate Reactions

• A type of double replacement

• Forms carbonic acid (H₂CO₃) which immediately breaks down into carbon dioxide and water

• Ex.

Note: remember to do this, or else points for a question like this will not count

4.6 - Combustion Reactions & Review of All Reaction Types

Combustion

• Reaction of oxygen with a compound containing carbon & hydrogen (called a hydrocarbon) which then produces carbon dioxide and water

• C_XH_Y + O₂ → CO_{2 (g)} + H₂O_(l)

• Commonly called burning, burning in the presence of oxygen

• exothermic reaction, produces heat/energy

• Since these are organic and don’t have many patterns, the phase symbol for the hydrogen will definitely be given

• Note: balance C or H to make the process easier & faster

Summary of All Reactions

• Synthesis

  • Combining of elements

• Decomposition

  • Breaking into elements

• Single replacement

  • Metal or non-metal replaces single element, single element must be more active

• Double replacement

  • Metals switch places, must produce a solid, liquid, or gas

• Acid base neutralization

  • Double replacement, acid + base → salt + water

• Acid carbonate 

  • Double replacement, carbonic acid deteriorates quickly, producing water and carbon dioxide

• Combustion

  • Hydrocarbon + oxygen → carbon dioxide + water + energy, exothermic, basically burning, energy created

4.7 - Dissociation Reactions, Formula Equations, & Net Ionic Equations

Types of equations

1. Formula equations (the ones above)

2. Complete ionic equations

3. Net ionic equations

Dissociation equations

• All ionic compounds & acids dissociate or  break up into ions when dissolved in water

• Polyatomic ions stay together

• Ex. 

Steps:

1.  Dissociate only the aqueous parts

2. Balance the # of atoms

3. Balance the charges of the atoms

Dissociation vs Decomposition

Complete ionic equation

• Dissociate anything aqueous

• Leave the (s), (l), (g)

• Put any bonds formed together

Net  ionic equation

• Cross out any spectator ions (any that are the same on both sides of the equation) from complete ionic equations

Ex. 

Formula Equation

Complete  Ionic Equation

Net Ionic Equation

Ex. 

Formula Equation

Complete  Ionic Equation

Net Ionic Equation

Ex. 

No reaction, everything cancels out, all are spectator ions

4.9 - Reaction Energy

• Particles involved in a reaction can either release or absorb energy to or from their surroundings

• Particles release energy when they form bonds & need to absorb energy in order to break them

• Energy is released when new chemical bonds are made in the products

• Endothermic if: energy supplied > energy released

• Exothermic if: energy released > energy supplied

Endothermic reactions

• Endo: into

• Energy supplied > energy released

• Heat (energy) taken in

• Temperature of the substance drops

• Products feel cold

• X + Y → XY + energy

Exothermic reactions

• Exo : exit

• Energy released > energy supplied

• Heat (energy) released

• Temperature of the substance rises

• Products feel hot

• XY + energy → X + Y