Chemistry

Question 1

1.18 understand how elements are arranged in the Periodic Table:

• in order of atomic number
• in groups and periods.

Answer 1

periods are going up in atomic number. The period determines how many shells it has

Elements in the same group have similar reactivities resulting from a common number of outer electrons. same group=similar properties

This allows chemists to make predictions about:

reactivity
physical properties
type of bonding

Question 2

1.19 understand how to deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table

Answer 2

Metals
conduct electricity
have oxides which are basic, reacting with acids to give a salt and water

Non - Metals
do not conduct electricity (except for graphite)
have oxides which are acidic or neutral

Question 3

1.20 understand how to use electrical conductivity and the acid-base character of oxides to classify elements as metals or non-metals

Answer 3

metals:
-conduct electricity because they allow charge to pass through them
-metal oxides are basic
this means they will neutralise acids
-metal oxides which dissolve will form solutions with a pH of more then 7
(pH is a measure of acidity)

Question 3

1.21 identify an element as a metal or a non-metal according to its position in the Periodic Table

Answer 3

-metals are on the left side
-non metals are on the right side
-they are separated by a zig zag going from boron to astatine

Question 3

1.22 understand how the electronic configuration of a main group element is related to its position in the Periodic Table

Answer 3

Number or numbers of circles = Period number
Number of electrons in outer shell = group number
Total number of electrons in all shells = Atomic number

Question 4

1.23 understand why elements in the same group of the Periodic Table have similar chemical properties

Answer 4

This is because they have the same amount of outer electrons so will react and bond similarly

Question 5

1.24 understand why the noble gases (Group 0) do not readily react

Answer 5

because they have a full outer shell of electrons. This means they’re not desperate to give up or gain electrons.
Atoms are most stable when they have full outer shells

Question 6

1.25 write word equations and balanced chemical equations (including state symbols):

• for reactions studied in this specification
• for unfamiliar reactions where suitable information is provided.

Answer 6

(g) means gas
(l) menas liquid
(s) means solid
(aq) means aqueous
Word equation: equation with the name of the reactants and the products rather than their formula
Balanced equation: equation with the formula of he reactants and the products and each element has to have an equal amount on each side of the equation

Question 7

1.26 calculate relative formula masses (including relative molecular masses) (Mr) from relative atomic masses (Ar)

Answer 7

Relative atomic mass (Mr): of a compound is the sum of the relative atomic masses of the atoms or ions in that compound
Relative atomic mass (Ar): of an element is the average mass of one atom of that element compared with 1/12 of the mass of a carbon-12 atom

Question 8

1.27 know that the mole (mol) is the unit for the amount of a substance

Answer 8

Symbol is mol
The mass of one mole of a substance in grams is numerically equal to its relative formula mass

Question 9

1.28 understand how to carry out calculations involving amount of substance, relative atomic mass (Ar) and relative formula mass (Mr)

Answer 9

moles = mass ÷ relative atomic mass
mass = moles x relative atomic mass

Question 10

1.29 calculate reacting masses using experimental data and chemical equations

Answer 10

Use equations in 1.28

Question 11

1.30 calculate percentage yield

Answer 11

Percentage yield = (actual mass of product made (from experiment) / maximum theoretical yield (from calculation)) x 100%
It is not always possible to obtain the calculated amount of product in an experiment because:
Not all of your reactants react
It didn’t go to completion
It wasn’t pure

Question 12

1.31 understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation

Answer 12

Example experiment to find formula of magnesium oxide:
weigh some pure magnesium
Heat magnesium to burning in a crucible to form magnesium oxide, as the magnesium will react with the oxygen in the air
weigh the mass of the magnesium oxide
Known quantities: mass of magnesium used & mass of magnesium oxide produced
Required calculations:
mass oxygen = mass magnesium oxide - mass magnesium
moles magnesium = mass magnesium ÷ molar mass magnesium
moles oxygen = mass oxygen ÷ molar mass oxygen www.pmt.ed
calculate ratio of moles of magnesium to moles of oxygen
use ratio to form empirical formula

Question 13

1.32 know what is meant by the terms empirical formula and molecular formula

Answer 13

Empirical formula: the simplest whole number ratio of elements in that compound
Molecular formula: the actual number of atoms in one molecule in that compound

Question 14

1.33 calculate empirical and molecular formulae from experimental data

Answer 14

Calculate the moles of each element
Divide the answers from step 1 by the smallest number to give a 1:X ratio
Find the nearest whole number ratio
Write the empirical formula

Question 15

1.36 practical: know how to determine the formula of a metal oxide by combustion (e.g. magnesium oxide) or by reduction (e.g. copper(II) oxide)

Answer 15

See 1.31

Question 16

2.1 understand how the similarities in the reactions of these elements with water provide evidence for their recognition as a family of elements

Answer 16

Group 1 Metals will React Similarly with Water as they are a Family of Elements called Alkali Metals
They will React Vigorously with Water to Produce an Alkaline Metal Hydroxide and Hydrogen Gas

Question 17

2.2 understand how the differences between the reactions of these elements with air and water provide evidence for the trend in reactivity in Group 1

Answer 17

The Reactivity of Group 1 Metals will Increase Down the Group
As You go Down the Group, the Observations during the Reaction with Air and Water becomes More Vigorous

Question 18

2.3 use knowledge of trends in Group 1 to predict the properties of other alkali metals

Answer 18

As the Reactivity of Alkali Metals Increases Down the Group, Rubidium, Caesium and Francium will React More Vigorously with Air and Water.
you would think the higher amount of electrons the more reactive it is for alkali metals

Question 18

2.5 know the colours, physical states (at room temperature) and trends in physical
properties of these elements

Answer 18

group 1:
gets lighter down the group
all solids
they get softer down the group
mass and density increases down the group
melting point decreases down the group
group 7:
gets darker down the group
gas to solid down the group
melting point increases down the group
group 0:
colourless
all gas
size and density increases down the group
melting point increases down the group

Question 19

2.6 use knowledge of trends in Group 7 to predict the properties of other halogens

Answer 19

Element Colour State at room temp
Fluorine (F2) Yellow Gas
Astatine (At2) Black Solid

Question 20

2.7 understand how displacement reactions involving halogens and halides provide
evidence for the trend in reactivity in Group 7

Answer 20

Halogens undergo redox reactions with metal halides in solution, displacing less reactive halogens from their compounds. These displacement reactions are used to establish an order of reactivity down Group 7 of the periodic table

Question 21

2.28 describe the use of litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions

Answer 21

Litmus, phenolphthalein, and methyl orange are indicators that change color in acidic and alkaline solutions, with litmus turning red in acids and blue in alkaline solutions, while phenolphthalein turns pink in alkaline solutions and methyl orange turns yellow in alkaline solutions.

Question 22

2.29 understand how to use the pH scale, from 0–14, can be used to classify solutions as strongly acidic (0–3), weakly acidic (4–6), neutral (7), weakly alkaline (8–10) and strongly alkaline (11–14)

Answer 22

The pH scale is like a color-changing indicator that helps classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline, or strongly alkaline, similar to how a universal indicator changes colors to identify the pH of a solution.

Question 23

2.30 describe the use of universal indicator to measure the approximate pH value of an aqueous solution

Answer 23

A universal indicator is a tool used to estimate the pH of a liquid by adding a few drops of indicator solution or dipping indicator paper into the liquid and matching the resulting colour change to a specific pH value on a scale.

Question 24

2.31 know that acids in aqueous solution are a source of hydrogen ions and alkalis in a aqueous solution are a source of hydroxide ions

Answer 24

Acids in water release hydrogen ions (H+), making the solution acidic, while alkalis in water release hydroxide ions (OH-), making the solution alkaline. Think of acids as sources of H+ and alkalis as sources of OH-.

Question 25

2.32 know that alkalis can neutralise acids

Answer 25

When an alkali (base) and an acid react, they create a neutral solution, forming water and a salt, such as chloride, nitrate, or sulfate, depending on the acid used.

Question 26

2.35 understand acids and bases in terms of proton transfer

Answer 26

Acids and bases are substances that can either lose or gain electrons, forming charged particles called ions, with acids losing electrons to form positive hydrogen ions (H+) and bases gaining electrons to form negative hydroxide ions (OH-).

Question 27

2.34 know the general rules for predicting the solubility of ionic compounds in water:

Answer 27

common sodium, potassium and ammonium compounds are soluble
all nitrates are soluble
common chlorides are soluble, except those of silver and lead(II)
common sulfates are soluble, except for those of barium, calcium and lead(II)
common carbonates are insoluble, except for those of sodium, potassium and
ammonium
common hydroxides are insoluble except for those of sodium, potassium and
calcium (calcium hydroxide is slightly soluble).

Question 28

2.36 understand that an acid is a proton donor and a base is a proton acceptor

Answer 28

Acids are like generous friends who willingly give away protons (H+) to make a solution acidic, while bases are like eager hosts who happily accept protons (H+) to make a solution alkaline.

Question 29

2.37 describe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals, bases and metal carbonates (excluding the reactions between nitric acid and metals)
to form salts

Answer 29

When hydrochloric acid, sulfuric acid, and nitric acid react with metals, bases, and metal carbonates, they form salts, just like when metal oxides, metal hydroxides, and ammonia react with acids to form salts and water.

Question 30

2.38 know that metal oxides, metal hydroxides and ammonia can act as bases, and that alkalis are bases that are soluble in water

Answer 30

Metal oxides, metal hydroxides, and ammonia are bases that can neutralize acids, and alkalis are bases that dissolve in water, which can be demonstrated by adding a small amount of insoluble base to an acid, filtering out the excess base, and allowing the remaining solution to dry and crystallize.

Question 31

2.39 describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant

Answer 31

1. Add some dilute acid to a beaker (warm acid in water bath)
2. Add a spatula of insoluble reactant and continue heating while mixing (mixture will effervesce) continue adding powder until unreacted powder is left over (add it in excess)
3. Filter of excess insoluble reactant and transfer the filtrate (solution) to an evaporating basin
   CRYSTALLISATION
4. Heat the solution over a Bunsen burner to boil pff so,me of the water and to concentrate the solution
5. Keep heating until a saturated solution is formed (yes: glass string rod in solutions if crystals form on tip = saturated)
6. Stop heating solution and allow it to cool slowly at room temp so larger crystals can form
7. Remove crystals from reaction mixture to filtration (was the crystals with distilled water to remove impurities
8. Crystals can be left to dry in a warm place

Question 32

2.42 practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting from copper(II) oxide

Answer 32

Add Dilute Sulfuric Acid into a beaker and heat using a bunsen burner flame
Add Copper (II) Oxide (insoluble base) in small amounts and stir until Copper (II) Oxide is in excess (stops disappearing)
Filter mixture into an evaporating basin to remove the excess Copper (II) Oxide
Leave filtrate in a warm place to dry and crystallize
Decant excess solution
Blot crystals dry

Question 33

1.1 understand the three states of matter in terms of the arrangement, movement and energy of the particles

Answer 33

Solids
Particles are close together and regularly packed
Particles vibrate around a fixed point
Particles have less kinetic energy than both liquid and gas

Liquids
Particles are close together but are irregular
Particles are free to move
Particles have less kinetic energy than gasses but more than solids

Gases
Particles are far apart and there are no forces between them
Particles are free to move
Particles have more kinetic energy than liquids and solids

Question 34

1.2 understand the interconversions between the three states of matter in terms of:
the names of the interconversions
how they are achieved
the changes in arrangement, movement and energy of the particles.

Answer 34

Melting = Solid -> liquid
Evaporating = liquid -> gas
Condensing = gas -> liquid
Freezing = liquid -> gas
Sublimation = Solid -> gas

Question 35

1.3 understand how the results of experiments involving the dilution of coloured solutions and diffusion of gases can be explained

Answer 35

Diffusion
-Movement of particles from an area of high concentration to an area of low concentration
-For this to work, particles must be able to move.Therefore diffusion does not occur in solid, since the particles cannot move just vibrate.
-Therefore, coloured solutions are diluted by adding water, because the particles of the colour diffuse to the air of low concentration, mixing with the water molecules, causing dilution to occur

Question 36

1.4 know what is meant by the terms:
* solvent
* solute
* solution
* saturated solution.

Answer 36

Solvent: Liquid in which a solute dissolves
Solute: substance that dissolves in a liquid to form a solution
solution: mixture formed when a solute has dissolved in a solvent
saturated solution: solution in which no more solvent can be dissolved

Question 37

1.14 know what is meant by the terms atom and molecule

Answer 37

-All element are made of atoms
-Element: All substances with one sort of atom
-An atom: The smallest part of an element
-Molecule: Formed when atoms join together by chemical bond

Question 38

1.15 know the structure of an atom in terms of the positions, relative masses and relative charges of sub-atomic particles

Answer 38

RELATIVE MASS
Proton: 1. Neutron: 1. Electron: 1/1836
RELATIVE CHARGE
Proton: +1. Neutron: 0. Electron: -1
POSITION
Proton: in the nucleus
Neutron: in the nucleus
Electron: in shells around the nucleus

Question 39

1.16 know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Ar)

Answer 39

Atomic number: Number of protons
Mass number: Number of protons + neutrons
Isotopes: Different atoms of the same element containing the same number of protons but different numbers of neutron in their nuclei

Question 40

1.17 be able to calculate the relative atomic mass of an element (Ar) from isotopic abundances

Answer 40

75% of chlorine atoms are the type 35Cl (have a mass number of 35)

25% of chlorine atoms are of the type 37Cl (have a mass number of 37)

In order to calculate the relative atomic mass (Ar) of chlorine, the following steps are used:

Multiply the mass of each isotope by its relative abundance
Add those together
Divide by the sum of the relative abundances (normally 100)
[ A_r = \frac{( (35 \times 75) + (37 \times 25) )}{100} ]

[ A_r = 35.5 ]

Question 41

1.37 understand how ions are formed by electron loss or gain

Answer 41

Ions are electrically charged particles formed when atoms lose or gain electrons. This loss or gain leaves a full outer shell, so the electronic structure of an ion is the same as that of a noble gas (such as helium, neon or argon).

Question 42

1.38 know the charges of these ions:

• metals in Groups 1, 2 and 3
• non-metals in Groups 5, 6 and 7
• Ag+, Cu2+, Fe2+, Fe3+, Pb2+, Zn2+
• hydrogen (H+), hydroxide (OH-), ammonium (NH4+), carbonate (CO32-),
  nitrate (NO3-), sulfate (SO42-).

Answer 42

Name of Ion
Sulfate SO42- -2
Carbonate CO32- -2
Nitrate NO3- -1
Hydroxide OH- -1
Ammonium NH4+ +1
Silver ion Ag+ +1
Zinc ion Zn2+ +2
Hydrogen ion H+ +1
Copper (II) ion Cu2+ +2
Iron (II) ion Fe2+ +2
Iron (III) ion Fe3+ +3
Lead (II) ion Pb2+ +2

Question 43

1.39 write formulae for compounds formed between the ions listed above

Answer 43

Writing the electron configuration of an atom allows you to work out the electron configuration of the ion and therefore the charge on the ion.

Examples:

Atom = Mg

Electron configuration = 2,8,2

remove the two electrons from the outer shell to achieve the same electron configuration as the nearest noble gas, Neon (Ne 2,8)

Ion = Mg2+

[2,8]
2+

Atom = O

Electron configuration = 2,6

add two electrons to the outer shell to achieve the same electron configuration as the nearest noble gas, Neon (Ne 2,8)

Ion = O2- [2,8]2-

Question 44

2.44 describe tests for these gases:

• hydrogen
• oxygen
• carbon dioxide
• ammonia
• chlorine

Answer 44

Gas Test Result if gas present
hydrogen (H2) Use a lit splint Gas pops
oxygen (O2) Use a glowing splint Glowing splint relights
carbon dioxide (CO2) Bubble the gas through limewater Limewater turns cloudy
ammonia (NH3) Use red litmus paper Turns damp red litmus paper blue
chlorine (Cl2) Use damp litmus paper Turns damp litmus paper white (bleaches)

Question 45

2.45 describe how to carry out a flame test

Answer 45

METHOD:

Platinum or Nichrome wire is cleaned by dipping it into Hydrochloric acid
End of wire is dipped into fresh Hydrochloric acid and then into solid sample
End of the wire with solid sample attached is placed into a non-roaring, non-luminous bunsen flame
Colour of flame is observed and recorded

Question 46

2.46 know the colours formed in flame tests for these cations:

• Li+ is red
• Na+ is yellow
• K+ is lilac
• Ca2+ is orange-red
• Cu2+ is blue-green.

Answer 46

Ion Colour in flame test
lithium (Li⁺) red
sodium (Na⁺) yellow
potassium (K⁺) lilac
calcium (Ca²⁺) orange-red
copper (II) (Cu²⁺) blue-green

Question 47

2.49 describe a test for the presence of water using anhydrous copper(II) sulfate

Answer 47

Add anhydrous copper (II) sulfate (CuSO4) to a sample.

If water is present the anhydrous copper (II) sulfate will change from white to blue.

Question 48

2.50 describe a physical test to show whether a sample of water is pure

Answer 48

If the sample is pure water it will boil at 100oC