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IB Chemistry HL > Chapter 1 > Flashcards

Chapter 1 Flashcards

Stoichiometric relationships (52 cards)

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1
Q

Stoichiometry

A

Relationship between amount of reactants and amount of products in a chemical reaction

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2
Q

Law of conservation of mass

A
  1. The sum of masses of all reactants must be equal to sum of masses of all products
  2. Matter cannot be created or destroyed in a chemical reaction
  3. No matter is gained or lost during chemical change
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3
Q

States of matter

A
  1. Solid
  2. Liquid
  3. Gas
  4. Plasma- ionised gas (mainly found in outer space)
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4
Q

Determination of the state of matter

A

Matter can exist in different states depending on temperature and pressure

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5
Q

Different physical properties of matter

A

Are due to:

  1. Different arrangement and movement of particles
  2. Depends on amount of kinetic energy of that particle
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6
Q

Kinetic energy

A

Energy related to the motion of an object

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7
Q

Properties of solids

A
  • particles are closely packed
  • strong forces between particles, vibrate about fixed positions
  • fixed shape
  • fixed volume
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8
Q

Properties of liquids

A
  • particles are more spread out than in solids
  • weaker forces between particles, they can move past each other
  • take the shape of the container
  • fixed volume
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9
Q

Properties of gases

A
  • particles are very spread out
  • negligible forces between particles, they move randomly
  • no fixed shape
  • no fixed volume
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10
Q

Changes of state

A

Directly related to changes in temperature

  • increase in temp. = increase in avg. KE of particles in a substance
  • when heated, particles gain KE + able to intermolecular forces that exist between them
  • results in a change of state
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11
Q

Intermolecular forces

A

Attractive forces that exist between particles

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12
Q

Changes of state:

  1. SOLID to LIQUID
  2. LIQUID to GAS
  3. SOLID to GAS (no liquid state)
  4. GAS to LIQUID
  5. LIQUID to SOLID
  6. GAS to SOLID (no liquid state)
A
  1. Melting (energy absorbed)
  2. Evaporation/Boiling (energy absorbed)
  3. Sublimation (energy absorbed)
  4. Condensation (energy released)
  5. Freezing (energy released)
  6. Deposition (energy released)
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13
Q

Evaporation vs. Boiling

A

Evaporation:

  • change of a liquid to a gas
  • BUT only takes place at surface of a liquid
  • can occur at temperatures below BP of liquid

Boiling:

  • change of a liquid to a gas
  • takes place throughout liquid (bubbles of gas are formed within the liquid, not only at surface)
  • occurs at a specific temperature
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14
Q

Steps of the heating curve of a solid

A
  1. Upward slopes, temp. increases as substance is heated
  2. KE of particles increases, and they vibrate faster
  3. BUT, during changes of state, temp. remains constant. At these points, energy is used to overcome intermolecular forces between particles
  4. During melting, energy input is used to overcome intermolecular forces that hold particles in solid in fixed positions
  5. During boiling, energy input is used to overcome intermolecular forces that hold particles in liquid together
  6. Hence, during changes of state, temp. remains constant
  7. Once change of state is complete, temp. starts to increase again
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15
Q

Elements

A
  • made up of same kind of atom
  • can’t be broken down by chemical means into a simpler substance
  • can be divided into metals, non-metals and metalloids (properties of both)
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16
Q

Atom

A
  • smallest particle that shows characteristic properties of that element
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17
Q

Compounds

A

A substance composed of two or more different elements that are chemically combined

  • when atoms of different elements react, they lose their characteristic properties
  • properties of the elements are replaced by those of new compound formed in chemical reaction
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18
Q

Molecule

A

Consists of two or more atoms that are chemically bonded together

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19
Q

Properties of molecules

A
  • not all molecules can be classes as compounds, as they are composed of same kind of atom bonded together
  • electrically neutral, although some have a dipole
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20
Q

Mixture

A

Composed of two or more substance that aren’t chemically combined- each substance retains its original properties

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21
Q

Homogeneous mixture

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A

Has uniform composition throughout eg. solutions

  • different components are NOT chemically combined
  • but don’t separate physically on standing
22
Q

Heterogeneous mixture

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A

Has a non-uniform composition throughout eg. air

- will separate into two layers on standing

23
Q

Solution

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A

Composed of a solute dissolved in a solvent

24
Q

Concentration of the solution

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A

Amount of solute dissolved in a known volume of a solution

25
Examples of solutions
1. Solid/solid: alloys eg. steel (iron and carbon) 2. Solid/liquid: copper sulphate solution 3. Liquid/liquid: wine (ethanol and water) 4. Gas/liquid: hydrochloric acid (hydrogen chloride in water)
26
Fractional distillation
Where a complex mixture can be separated into fractions because the different components have different BP - fractions with lower BP rise higher up the column, where it's cooler before they condense
27
Chemical reactions
Involve formation of new chemical substances - differ from physical changes that only involve a change of state with no new substances being formed - atoms present in reactants are rearranged and join together to form products - chemical change taking place during a chemical reaction can be represented by a chemical equation
28
Law of conservation of mass
Mass is conserved in a chemical reaction - number and type of each atom must be the same in the reactants and products - in a chemical reaction, atoms aren't created or destroyed, they are only rearranged to form new products
29
Diatomic molecules
Two atoms bonded together | - 7 elements that are diatomic (Hydrogen, oxygen, nitrogen) and then group 17 elements
30
Relative atomic mass (Ar)
The weighted average mass of an atom compared to 1/12 mass of an atom of Carbon-12 - allows comparison of the relative masses of different atoms
31
Atomic mass unit
Mass of 1/12 of an atom of Carbon-12 | - atom of C-12 has a mass of exactly 12 amu
32
Why is Carbon-12 used as the reference for RAM?
Its mass can be accurately measured | - most abundant stable isotope of carbon
33
Natural sample of an element
- made up of a mixture of isotopes according to their relative abundance
34
Isotope
An atom of the same element with same number of protons, different number of neutrons
35
Mass spectrometer
- Abundance of isotopes of an element can be analysed using a mass spectrometer
36
Relative formula mass and relative molecular mass (Mr)
Weighed average mass of the compound compared to 1/12 the mass of an atom of Carbon-12 - allows comparison of relative masses of different compounds
37
RFM vs. RMM | Relative formula mass vs. Relative molecular mass
RFM: sum of the RAM of all atoms in a formula unit of a substance - mainly used for ionic compounds, that don't form molecules RMM: sum of the RAM of all atoms in a formula unit of a substance - used for molecular covalent substances
38
Mole concept
Used to count very small particles eg. atoms, molecules and ions - mole concept is necessary because individual atoms are so small, it's impossible to count them
39
Mole
Amount of a substance (n) that contains the same no. of particles as there are atoms in exactly 12g of isotope C-12
40
Avogadro's constant
- no. of particles in a mole of a substance | - value is 6.02 x 10^23 (no. of atoms in exactly 12g of isotope C-12)
41
Properties of moles
- a mole of any substance contains same no. of particles, although moles of different substances have different masses
42
Molar mass (M)
Mass in grams of one mole of a substance Unit: g/mol Molar mass of a substance is equal to its RAM (RFM/RMM)
43
What is a particle?
Particle = atoms, molecules, ions, formula units and electrons In one mole of a substance there are 6.02 x 10^23 particles
44
Equation for moles, mass and molar mass
n (amount in mol) m (mass in grams) M (molar mass) n = m/M
45
Equation for moles and no. of particles
n = no. of particles/ 6.02 x 12^23
46
Equation for the mass of one molecule
Mass of one molecule = Molar mass / 6.02 x 10^23
47
Empirical formula
Simplest whole number ratio of atoms in a compound
48
Molecular formula
Actual number of atoms in a compound
49
Calculating empirical formula
1. From percentage, find mass of each element in 100g of compound (% given = mass in g) 2. Find amount in mol of each element present 3. Divide each element by lowest value obtained 4. Empirical formula ratio- find the simplest whole-number ratio
50
Percentage composition by mass
Expresses the mass of each element in a compound as a percentage Equation: % composition = (mass of element in compound/ molar mass of compound) x 100
51
Experimental methods of determining empirical formula
1. Heat the substance, measure changes in mass 2. Calculate water of crystallisation of a hydrated compound 3. Burning a sample of a substance in an excess of oxygen (complete combustion) and analysing products of reaction
52
Water of crystallisation
Determined experimentally by heating a hydrated salt until the water evaporates, leaving the anhydrous salt mass of hydrous salt - mass of anhydrous salt = mass of water evaporated

IB Chemistry HL (9 decks)

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