Grade 11 HL Chemistry: Unit 1

Accuracy

How closely value is to true or accepted (theoretical) value

Precision

How close successive measurements are to each other

More decimal places = more precise

Systematic error

Consistent and repeatable results in values too low or too high, affects accuracy

Addressed through procedure change/using same equipment

Random error

Increasing spread, error without pattern, affects precision

Addressed by taking more measurements/more precise instruments

Uncertainty

Absolute (red, decimal) + relative (%)

Average uncertainty

ΣTrials uncertainties/# of trials

Multiplying/Dividing Uncertainties

Sum of relative then convert to absolute

Analog

Lines/ticks

Significant digits: Digits for sure + one additional estimate

Uncertainty: Half of smallest tick (Ex. lower is mm → ±0.05cm)

Significant digits rules

Non-zero digits are significant

Leading zeroes are not significant.

Zeros in-between non-zero digits are significant

Zeros to the right of a non-zero is only significant if it is also to the right of a decimal place

Counted values: Certain of amount (Ex. # of apples) → infinitely significant

Constants: Values determined by someone (Ex. π) → infinitely significant

Numbers from formula: Math formula (C = 2π) → infinitely significant

Odd-up rounding rule

When number = 5 if odd → up, even → down

Multiplying/Dividing sig figs

Same as value with least number of significant figures

Addition/Subtraction sig figs

Same as value with least number of decimal places

Model

Explains unfamiliar facts in nature but can be changed with new evidence

Aristotle/Plato

dudes believing in the 4 elements

Atomos

Not to be cut

Dalton's atomic theory:

All matter is made of indivisible particles called atoms

Atoms of all the same elements are similar in slope/mass

Atoms of different elements differ in slope/mass

Atoms can neither be created nor destroyed

Atoms combine in whole number ratios to form compounds

Thomson's plum pudding model

Electricity is passed through a gas at one end of the tube, and an invisible beam of particles "cathode rays" travels through the tube and hits a fluorescent material that glows. Stream bent towards the positive plate and away from the negative

Called them "corpuscles" today known as electrons

Atoms are neutral overall, so the rest of the atom is positively charged

Rutherford's Gold Foil Experiment

discovered the nucleus

Bohr Model

Argued that previous model was incorrect because negative electrons moving randomly around a nucleus would become attracted to the positive nucleus, gradually lose their energy, and collapse into the nucleus

Experiment with heating different elements. Each element emitted different colours of light after it's heated. Splitting the light produced by a heated element using a prism produced a different spectrum for each element.

When an electron moves between energy levels it either absorbs or releases a specific amount of energy. Each colour of light corresponds with a specific amount of energy that is released.

Electron mass

1/2000 relative mass

Standard atomic notation

Periodic table groups

Group 1: Alkali Metal

Group 2 : Alkaline Earth Metal

Transition Metals

Rare Earth Metal

• Lanthanide and Actinide Series, pulled out of the table

Group 17: Halogen

Group 18: Inert/Noble Gas

Weighted Average Atomic Mass

Atomic mass → in relation to carbon (Ex. hydrogen = 1/12 carbon)

Radioisotopes

Isotope that has an unstable nucleus, radioactive and can spontaneously decay

Nucleus breaks down, energy is released as radiation

Factors that affect trends

Nuclear charge/Nuclear force

Number of energy levels

Electron shielding: Electrons blocked by inner electrons

Electron-electron repulsion: Electrons in same shell repel each other but there are more protons, so protons overpowers → generally only for ionic radius

Atomic radius

Distance from nucleus to valence

Ionic radius

Cations smaller, anions bigger (EE repulsion)

Electron affinity

Energy required/released to gain one external electron to form an ion

Affinity = to like → easier to add electron = higher affinity

First ionization energy

Energy required/released to remove one valence electron to form an cation

Harder to remove electron = higher first ionization energy

Electronegativity

An atom's ability to attract shared electrons in a covalent bond to itself

Stronger attraction = greater electronegativity

Helps determine type of bond

• Pure covalent, Polar covalent bond, Ionic bond

Pure covalent

Two atoms with the same electronegativities

Polar covalent bond

When electrons are shared unequally, different electronegativities

Metallic behavior

Don't hold on to electrons tightly = more metallic properties

Melting point of metals

Temperature remains same while changing states

Particles move so fast the move away from each other, breaking force between particles

Greater attractive force of nucleus to valence electrons = greater melting point

Increases towards the right and up

Reactivity

Relative ease at which an atom will gain or lose an electron to become an ion

Percent of error

Improvements Dalton made

Atoms can neither be created nor destroyed

Atoms combine in whole number ratios to form compounds

What observations did Dalton use to create his model?

Took apart/made compounds from elements

Mass measurement mass of products = mass of reactants

Elements combined in definite ratios

Ionic bonding

Electrostatic force of attraction between a cation and anion

THE FORCE NOT THE TRANSFER!!!

3D crystalline structure formed by repeating ions

Oppositely charged ions next to each other

Formula unit

Smallest reduced ratio of ions in a 3D crystal

What affects strength of ionic bonds?

Smaller ions pack closer and attract each other with greater force

Magnitude of charge: Higher charge = Higher ionic bond strength

Covalent bonding

Creates molecular compound

Electrostatic attraction between the nuclei of 2 atoms and a shared pair of valence electrons

Lewis diagram can't form a ring

Covalent bond

Shared pair of electrons

Metallic Bond

Held together by the sharing of electrons between metal atoms

Coordination covalent bond

Both electrons from the shared pair originate from one atom

Formal charge

Group number - covalent bonds - electron lone pairs

Resonance structure

Two+ Lewis dot structures describe a molecule equally well

Simply (binary) ionic compound nomenclature

Name: Cation + anion + ide

Reduce to simplest ratio

Multivalent metals

Metals that can form an ion in more than one way, resulting in ions with different charges

Antimony (Sb)

Chromium (Cr)

Cobalt (Co)

Copper (Cu)