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OCR A Chemistry > Chemistry__Module 2 > Flashcards

Chemistry__Module 2 Flashcards

1
Q

What is the charge of a proton?

A

+1

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2
Q

What is the mass of a proton?

A

1

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3
Q

What is the charge of a neutron?

A

0

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4
Q

What is the mass of a neutron?

A

1

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5
Q

What is the charge of an electron?

A

-1

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6
Q

What is the mass of an electron?

A

1/2000

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7
Q

What is the atomic number?

A

The number of protons in the nucleus of an atom.

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8
Q

What is the mass number?

A

Sum of protons and neutrons in the nucleus.

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9
Q

How do you work out the number of neutrons?

A

Mass number - atomic number.

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10
Q

What are ions?

A

Charged particles that are formed when an atom loses or gains electrons.

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11
Q

What charge do elements have if they gain electrons?

A

Negative charge, more electrons than protons.

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12
Q

What charge is gained when electrons are lost?

A

Positive charge, fewer electrons than protons.

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13
Q

What are isotopes?

A

Atoms of the same element with the same number of protons but a different number of neutrons.

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14
Q

Why do different isotopes of the same element react in the same way?

A

Number and arrangement of electrons decide the chemical properties; isotopes have the same configuration of electrons, so they have the same chemical properties; neutrons have no impact on reactivity.

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15
Q

What is relative atomic mass?

A

Weighted mean mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12.

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16
Q

What is relative isotopic mass?

A

The mass of an atom of an isotope compared with 1/12th of the mass of an atom of carbon-12.

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17
Q

How to calculate isotopic abundances?

A

Multiply each isotope by abundance, add up results, and divide by 100.

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18
Q

What are the uses of mass spectrometry?

A

To identify unknown compounds, find relative abundance of each isotope of an element, and determine structural information.

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19
Q

How is mass spectra represented?

A

Image representation.

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20
Q

How to calculate relative atomic mass from mass spectra?

A

Multiply each relative isotopic mass by its relative isotopic abundance, add results, and divide by the sum of the isotopic abundances.

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21
Q

What are positively charged ions called?

A

Cations.

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22
Q

What are negatively charged ions called?

A

Anions.

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23
Q

What is the charge of a nitrate ion?

A

NO3-

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24
Q

What is the charge of carbonate?

A

CO32-

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25
Q

What is the charge of sulfate?

A

SO42-

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26
Q

What is the charge of hydroxide?

A

OH-

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27
Q

What is the charge of ammonium?

A

NH4+

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28
Q

What is the charge of zinc ion?

A

Zn2+

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29
Q

What is the charge of silver ion?

A

Ag+

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30
Q

What are molecular ions?

A

Covalently bonded atoms that lose or gain electrons.

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31
Q

What is an empirical formula?

A

The simplest whole number ratio of atoms of each element present in a compound.

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32
Q

What is a molecular formula?

A

The number and type of atoms of each element of a molecule.

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33
Q

How to calculate empirical formula?

A

Divide the amount of each element by Mr, divide answers by the smallest, and if there’s a decimal divide/round to whole number.

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34
Q

What is relative molecular mass?

A

Average mass of a molecule compared to 1/12th of the mass of one atom of carbon-12.

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35
Q

What is a mole?

A

The amount of substance in grams that has the same number of particles as there are atoms in 12 grams of carbon-12.

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36
Q

What is the formula for moles?

A

Moles = mass/molar mass.

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37
Q

What is molar mass?

A

Mass in grams of 1 mole of a substance gmol-1.

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38
Q

What is Avogadro’s constant?

A

The number of atoms per mole of the carbon-12 isotope: 6.02 x 1023.

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39
Q

What is Avogadro’s law?

A

Under the same temperature and pressure, one mole of any gas would occupy the same volume.

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40
Q

What is molar gas volume?

A

Volume of one mole of a gas at a given temperature and pressure.

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41
Q

What is the molar gas volume at room temperature and pressure?

A

24000 cm3/24 dm3.

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42
Q

What is the formula linking moles, volume, and molar gas volume?

A

n = volume(dm3)/molar gas volume (24 dm3).

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43
Q

What is the ideal gas equation?

A

Pv = nRT, where P = pressure (Pa), V = volume (m3), n = number of moles (mol), R = gas constant (8.314 Jk-1 mol-1), T = temp (K).

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44
Q

How do you convert 1 cm3 to m3?

A

1 x 10-6 m3.

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45
Q

How do you convert 1 dm3 to m3?

A

1 x 10-3 m3.

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46
Q

How do you convert °C to K?

A

°C + 273.

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47
Q

What is 1 kPa in Pa?

A

1000 Pa.

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48
Q

What is 1 atm in Pa?

A

101325 Pa.

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49
Q

What is the formula linking moles, concentration, and volume?

A

c = n/v, where c = moldm-1, n = mol, v = dm3.

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50
Q

What are the 4 state symbols?

A

(s) solid, (aq) aqueous, (l) liquid, (g) gas.

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51
Q

What is an atomic orbital?

A

Region around a nucleus that can hold up to two electrons with opposite spins.

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52
Q

What are the names of the 4 subshells?

A

S, P, D, F.

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53
Q

How many orbitals does the S subshell have?

A

1 orbital.

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54
Q

How many electrons can the S subshell hold?

A

2 electrons.

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55
Q

How many orbitals does the P subshell have?

A

3 orbitals.

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56
Q

How many electrons can the P subshell hold?

A

6 electrons (3x2).

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57
Q

How many orbitals does the D subshell have?

A

5 orbitals.

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58
Q

How many electrons can the D subshell hold?

A

10 electrons (5x2).

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59
Q

How many orbitals does the F subshell have?

A

7 orbitals.

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60
Q

How many electrons can the F subshell hold?

A

14 electrons (7x2).

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61
Q

What does the principal quantum number show?

A

The shell occupied by the electrons.

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62
Q

What is a shell?

A

A group of orbitals with the same principal quantum number.

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63
Q

How many electrons can the first shell hold?

A

2.

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64
Q

How many electrons can the second shell hold?

A

8.

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65
Q

How many electrons can the third shell hold?

A

18.

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66
Q

How many electrons can the fourth shell hold?

A

32.

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67
Q

What are the shapes of the S and P orbitals?

A

Spherical and dumbbells.

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68
Q

For principal level 1, what is the subshell?

A

1s.

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69
Q

For principal level 2, what are the subshells?

A

2s, 2p.

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70
Q

For principal level 3, what are the sub levels?

A

3s, 3p, 3d.

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71
Q

For principal level 4, what are the sub levels?

A

4s, 4p, 4d, 4f.

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72
Q

What are the rules of electron configuration?

A

Added one at a time, lowest energy level is filled first, each energy level must be filled before the next one can fill, each orbital is filled single before pairing, and 4s is filled before 3d.

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73
Q

Why does 4s fill before 3d?

A

4s has a lower energy level than 3d before it is filled.

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74
Q

What is the electron configuration order?

A

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6.

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75
Q

What electrons are lost when an atom becomes a positive ion?

A

Electrons in the highest energy levels.

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76
Q

How do we determine the mass number and abundance of isotopes?

A

Use a mass spectrometer.

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77
Q

What is ionic bonding?

A

Electrostatic attraction between positive and negative ions.

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78
Q

What is the charge on a hydroxide ion?

A

OH-

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79
Q

What is the charge on a sulfate ion?

A

SO42-

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80
Q

What is the charge on a nitrate ion?

A

NO3-

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81
Q

What is the charge on a carbonate ion?

A

CO32-

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82
Q

What is the charge on an ammonium ion?

A

NH4+

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83
Q

What is the structure of an ionic compound?

A

Giant ionic lattice.

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84
Q

Why is an ionic compound structure a giant ionic lattice?

A

Because of the oppositely charged ions strongly attracted in all directions.

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85
Q

Why do ionic compounds have a high melting/boiling point?

A

Many strong electrostatic forces of attraction; a lot of energy is needed to overcome them.

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86
Q

Why do ionic compounds conduct electricity when molten/dissolved?

A

In a solution, the ions are free to move and can carry a charge.

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87
Q

Why do ionic compounds dissolve?

A

They have positive/negative ions, water is polar, so positive and negative ions are attracted to the slightly negative/positive parts of water, breaking up the structure.

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88
Q

What is covalent bonding?

A

Strong electrostatic attraction between a charged pair of electrons and the nuclei of the bonded atoms.

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89
Q

What is dative/co-ordinate covalent bonding?

A

Where one atom donates 2 electrons to form an atom/ion to form a bond.

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90
Q

What is average bond enthalpy?

A

Used as a measurement of covalent bond strength; the larger the value, the stronger the covalent bond.

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91
Q

What is the structure of a covalent bond?

A

Simple molecular.

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92
Q

What bonds are in simple covalent molecules?

A

Intermolecular forces, induced dipole-dipole, permanent dipole-dipole, hydrogen bonds.

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93
Q

Why do simple molecular forces have low boiling and melting points?

A

Weak intermolecular forces between molecules.

94
Q

Why can’t simple molecular forces conduct electricity?

A

They contain no ions.

95
Q

What is the structure of a covalent bond?

A

Simple molecular

96
Q

What bonds are in simple covalent molecules?

A

Intermolecular forces, induced dipole-dipole, permanent dipole-dipole, hydrogen bonds

97
Q

Why do simple molecular forces have low boiling and melting points?

A

Weak intermolecular forces between molecules

98
Q

Why can’t simple molecular forces conduct electricity?

A

They contain no ions

99
Q

What is metallic bonding?

A

Electrostatic attraction between positive metal ions and a sea of delocalised electrons

100
Q

What is a lone pair?

A

Electrons in the outer shell that are not involved in bonding

101
Q

Why do molecules have a specific shape with specific angles?

A

Because bonds repel each other equally

102
Q

What repels the most?

A

Lone pair + lone pair

103
Q

What repels the second biggest?

A

Lone pair + bond pair

104
Q

What repels the smallest?

A

Bond pair and bond pair

105
Q

What is the shape and bond angle for 2 bond pairs and 0 lone pairs?

A

Linear, 180

106
Q

Examples of linear molecules?

A

CO2, CS2, HCN, BeF2

107
Q

What is the shape and bond angle for 3 bond pairs and 0 lone pairs?

A

Trigonal planar, 120

108
Q

Examples of trigonal planar molecules?

A

BF3, AlCl3, SO3, NO3-, CO3^2-

109
Q

What is the shape and bond angle for 4 bond pairs and 0 lone pairs?

A

Tetrahedral, 109.5

110
Q

Examples of tetrahedral molecules?

A

CH3, SiCl4, SO4^2-, ClO4, NH4+

111
Q

What is the shape and bond angle for 5 bond pairs and 0 lone pairs?

A

Trigonal bipyramidal, 120 and 90

112
Q

What is the shape and bond angle for 2 bond pairs and 2 lone pairs?

113
Q

Examples of bent molecules?

A

H2O, OCl2, H2S, OF2, SCl2

114
Q

What is the shape and bond angle for 3 bond pairs and 2 lone pairs?

A

Trigonal planar, 120

115
Q

What is the shape and bond angle for 4 bond pairs and 2 lone pairs?

A

Square planar, 90

116
Q

What is the shape and bond angle for 6 bond pairs and 0 lone pairs?

A

Octahedral, 90

117
Q

What is the shape and bond angle for 3 bond pairs and 1 lone pair?

A

Trigonal pyramidal, 107

118
Q

Examples of trigonal pyramidal molecules?

A

NH3, NCl3, PF3, ClO3

119
Q

What is the VSEPR theory to find the shape of a molecule?

A
  1. Find the central atom. 2. Find valence/outer shell electrons. 3. Add one electron for each bonding pair. 4. Add/subtract electrons for the charge. 5. Positive charge then subtract. 6. Negative charge then add. 7. Add/divide by 2 to find electron pairs.
120
Q

What are induced dipole-dipole interactions?

A

They occur between all molecular substances and noble gases and are also known as London forces.

121
Q

What happens to the induced dipole-dipole forces as the molecule/atom gets bigger?

A

The more induced dipole-dipole forces and larger the electron clouds.

122
Q

How does the number of electrons affect the induced dipole-dipole interactions?

A

More electrons lead to a higher chance that temporary dipoles will form, making induced dipole-dipole interactions stronger between the molecules, so boiling points will be greater.

123
Q

How does the shape of the molecule affect the size of the induced dipole-dipole forces?

A

Long chain alkenes have a larger surface area of contact between the molecules for the interactions to form.

124
Q

What are permanent dipole-dipole interactions?

A

Weak electrostatic forces that exist between molecules with polarity.

125
Q

What is hydrogen bonding?

A

Occurs when hydrogen on one molecule forms a bond with the lone pair on nitrogen, oxygen, and fluorine. It has permanent and induced dipole-dipole forces.

126
Q

Why is ice less dense than water?

A

Hydrogen bonding keeps the molecules far apart.

127
Q

Why does water have a high melting/boiling point?

A

Hydrogen bonds are stronger than other intermolecular forces, so extra strength is required to overcome them.

128
Q

What happens to the boiling point as you go down group 7?

A

It increases; the number of electrons increases and the strength of London forces also increases.

129
Q

What is electronegativity?

A

The ability of an atom to attract a pair of electrons towards itself in a covalent bond.

130
Q

What is the most electronegative element?

131
Q

What does a bigger difference in electronegativity mean?

A

More ionic the compound will be.

132
Q

How can a bond become polar?

A

If the atoms attached have a difference in electronegativity; the bigger the difference, the more polar it will be.

133
Q

What are the intermolecular forces in terms of their strength?

A
  1. Hydrogen bonding 2. Permanent dipole-dipole 3. Induced dipole-dipole
134
Q

How do you convert cm³ to dm³?

A

Divide by 1000

135
Q

How do you convert dm³ to m³?

A

Divide by 1000

136
Q

What is percentage yield?

A

The efficiency of which reactants are converted into products.

137
Q

What is the percentage yield formula?

A

Actual yield in mol / theoretical yield in mol x 100

138
Q

Why might percentage yield not be 100%?

A
  1. Reaction may not be at equilibrium and not go to completion. 2. Reactants may be impure. 3. Side reactions could happen leading to by-products. 4. Reactants/products left behind in the apparatus. 5. Loss of products due to purification.
139
Q

What is atom economy?

A

A measure of the proportion of desired products compared with all the products formed in the reaction.

140
Q

What is the atom economy formula?

A

Molecular mass of desired products / sum of all the molecular masses of all the products x 100

141
Q

Why does a 100% percentage yield not mean 100% atom economy?

A

Even if all the reactants are converted into products, not all of the products of the reaction will be required.

142
Q

What reactions have 100% atom economy?

A

Addition reactions where the reactants combine to form a single product.

143
Q

What are substitution reactions?

A

Where the atoms from one reactant are swapped with another reactant, resulting in at least one product.

144
Q

What are the costs of low atom economy?

A
  1. A lot of waste and costs money to separate the desired products from the waste products. 2. Reactant chemicals are expensive. 3. Raw materials are limited. 4. Reduction conditions cost a lot to maintain.
145
Q

What are acids?

A

Proton donors.

146
Q

What do acids release?

A

H+ ions in aqueous solutions.

147
Q

What are strong acids?

A

Completely dissociate in water, releasing all the H+ ions.

148
Q

What are weak acids?

A

Partially dissociate in water.

149
Q

Examples of a weak acid?

A

CH3COOH (ethanoic acid) where the backward reaction is favored, so not many H+ ions are produced.

150
Q

Examples of strong acids?

A

HCl, H2SO4, HNO3

151
Q

Why do strong acids release lots of H+ ions?

A

The forward reaction is favored, so lots of H+ are produced.

152
Q

What are bases?

A

Proton acceptors.

153
Q

Examples of strong bases?

154
Q

Why are strong bases stronger than weak bases?

A

In strong bases, the forward reaction is favored strongly, so lots of OH- ions are produced. In weak bases, the backward reaction is favored, so not as many OH- ions are produced.

155
Q

Example of a weak base?

A

Ammonium (NH3)

156
Q

What are alkalis?

A

Bases that can dissolve in water to form aqueous hydroxide ions.

157
Q

What are polyprotic acids?

A

When acids can donate more than one proton.

158
Q

What type of acid is HNO3?

A

Monoprotic; 1 mol will produce 1 mol of H+ ions.

159
Q

What type of acid is H2SO4?

A

Diprotic; 1 mol will produce 2 mols of H+ ions.

160
Q

What type of acid is H3PO4?

A

Triprotic; 1 mol will produce 3 mols of H+ ions.

161
Q

Why do acids react with bases in neutralization reactions?

A

H+ ions produced from acids react with OH- ions produced from alkalis to form salt and water. H+ + OH- → H2O.

162
Q

Why is ammonia an exception from other bases?

A

Ammonium doesn’t produce OH- ions directly and reacts with water first, accepting a proton to produce ammonium ions and OH-. NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq).

163
Q

What happens when a metal reacts with acids?

A

Metal + acid → salt + hydrogen.

164
Q

What happens when metal oxides react with acids?

A

Metal oxide + acid → salt + water.

165
Q

What happens when metal hydroxide reacts with acid?

A

Metal hydroxide + acid → salt + water.

166
Q

What happens when metal carbonate reacts with acid?

A

Salt + carbon dioxide + water.

167
Q

What is a salt?

A

A compound formed when H+ of an acid is replaced by a metal ion or positive ion.

168
Q

How are ammonium salts formed?

A

Ammonium + acid → ammonium salt.

169
Q

What is the charge on the phosphate ion?

170
Q

What is the charge on the carbonate ion?

171
Q

What is the charge on the sulfate ion?

172
Q

What is the charge on the nitrate ion?

173
Q

What is the charge on the hydroxide ion?

174
Q

What is the charge on the ammonium ion?

175
Q

What are hydrated crystals?

A

Crystalline structures containing water molecules.

176
Q

What are anhydrous crystals?

A

Crystalline structures that contain no water molecules.

177
Q

What is water of crystallization?

A

Water molecules that form an essential part of the crystalline structure of a compound.

178
Q

How to calculate the water of crystallization?

A
  1. Calculate the number of moles. 2. Divide by the smallest.
179
Q

Why must the sample containing isotopes be ionized?

A
  1. To accelerate the process. 2. To deflect.
180
Q

What can be adjusted in the mass spectrometer to enable ions formed by the different isotopes to be directed into the detector?

A

Electric field.

181
Q

How do van der Waals forces arise?

A
  1. Uneven distribution of electrons. 2. Creates temporary dipoles in molecules. 3. Causes induced dipoles in neighboring molecules.
182
Q

How do you work out total relative atomic mass?

A

Isotopic mass x abundance / total abundance. Total abundance is 100 if the abundance is in percentage.

183
Q

How do you work out relative atomic mass if the relative abundance is used instead of percentage abundance?

A

Isotopic mass x relative abundance / total relative abundance.

184
Q

What are the conversions in chemistry?

A

1000 mg = 1 g, 1000 g = 1 kg, 1000 kg = 1 tonne.

185
Q

What is the oxidation number for group 1 metals?

186
Q

What is the oxidation number for group 2 metals?

187
Q

What is total abundance?

A

Total abundance is 100 if the abundance is in percentage.

188
Q

How do you work out relative atomic mass using relative abundance?

A

Isotopic mass x relative abundance / total relative abundance.

189
Q

What are the conversions in chemistry?

A

1000 mg = 1 g
1000 g = 1 kg
1000 kg = 1 tonne.

190
Q

What is the oxidation number for Al?

191
Q

What is the oxidation number for H?

A

+1 except in hydrides where it’s -1 like in NaH.

192
Q

What is the oxidation number of F?

193
Q

What is the oxidation number of Cl, Br, and I?

A

-1 except in compounds with oxygen and fluorine.

194
Q

What is the oxidation number of O?

A

-2

Except in peroxides (H₂O₂) = -1; Compounds with Fluorine = -1.

195
Q

What is the method for heating in a crucible?

A
  1. Weigh an empty clean dry crucible and lid.
  2. Add the hydrated salt to the crucible and weigh again.
  3. Heat strongly with a Bunsen burner for a couple of minutes.
  4. Weigh the contents again.
  5. Heat crucible until you reach constant mass to ensure reaction is complete.
196
Q

What is the method for making a standard solution?

A
  1. Weigh the solid using a balance and weighing boat to 2 dp.
  2. Transfer to beaker and re-weigh the weighing boat to record difference in mass.
  3. Wash the solid left over using de-ionised water.
  4. Add more de-ionised water until solid dissolved fully and stir with glass rod.
  5. Transfer to volumetric flask (250 cm³) using funnel, rinse beaker, glass rod, and funnel.
  6. Fill volumetric flask until graduation line and use pipette until you get to the bottom of meniscus.
  7. Invert flask with lid to ensure solution thoroughly mixed.
197
Q

Why do we use deionised water when making a standard solution?

A

So no ions react.

198
Q

Why is using a volumetric pipette more accurate than a measuring cylinder?

A

Because it has a smaller uncertainty.

199
Q

What does diluting a solution mean in terms of moles?

A
  1. Doesn’t change the amount of moles.
  2. Increases the volume.
  3. Lowers concentration.
200
Q

How to calculate new diluted concentration?

A

Refer to the provided image.

201
Q

What are irritants and how to be safe?

A

Dilute acids and alkalis.

Wear goggles.

202
Q

What are corrosives and how to be safe?

A

Stronger acids and alkalis.

Wear goggles.

203
Q

What does flammable mean?

A

Keep away from naked flame.

204
Q

What does toxic mean and how can you be safe?

A
  1. Wear gloves.
  2. Avoid skin contact.
  3. Wash hands after use.
205
Q

What does oxidising mean?

A

Keep away from flammable/easily oxidised materials.

206
Q

What goes in the burette?

A

Acid/alkali with known concentration.

207
Q

What goes in the conical flask?

A

Unknown concentration of acid/alkali and a few drops of indicator.

208
Q

What is the method for carrying out titrations?

A
  1. Pipette 25 cm³ of unknown alkali into conical flask.
  2. Add known acid solution to burette.
  3. Add indicator into conical flask.
  4. Use a white tile underneath to help observe the colour change.
  5. Add acid to alkali whilst swirling.
  6. Add the acid dropwise at the end point.
  7. Note burette reading before and after addition of acid.
  8. Repeat titration until 2 concordant results are obtained.
209
Q

How does phenolphthalein indicator work?

A

Pink - alkali.
Colourless - acid.
End point - pink colour disappears.

210
Q

How does methyl orange indicator work?

A

Yellow - alkali.
Red - acid.
End point is orange.

211
Q

What are concordant results?

A

Results within 0.10 cm³ of each other.

212
Q

Why do we only add a few drops of indicator?

A

Because they are generally weak acids, so if we add too much, then it will affect the titration results.

213
Q

What is the uncertainty of a balance?

A

+/- 0.001 g (3 dp balance).

214
Q

What is the uncertainty of a volumetric flask?

A

+/- 0.1 cm³.

215
Q

What is the uncertainty of a 25 cm³ pipette?

A

+/- 0.1 cm³.

216
Q

What is the uncertainty of a burette (start and end readings)?

A

+/- 0.10 cm³.

217
Q

How to calculate percentage uncertainty?

A

Refer to the provided image.

218
Q

How can you reduce uncertainties in a titration?

A

Replacing measuring cylinders with pipettes or burettes which have lower apparatus uncertainty lowers the percentage uncertainty.

219
Q

How can you reduce percentage uncertainty in a burette?

A
  1. Make the titre a larger volume by:
    • Increasing the volume and concentration of the substance in conical flask.
    • Decreasing the concentration of substance in the burette.
220
Q

How can you reduce uncertainty in measuring mass?

A
  1. Using a balance that measures to more decimal places.
  2. Using a larger mass.
221
Q

How to calculate percentage difference?

A

Real value - measured value x 100.

222
Q

What does it mean if the percentage uncertainty due to apparatus is less than percentage difference between actual value and calculated value?

A

There’s discrepancy in the result due to other errors.

223
Q

What does it mean if the percentage uncertainty due to apparatus is less than the percentage difference between the actual value and calculated value?

A

There’s no discrepancy and all errors in the results can be explained by the sensitivity of the equipment.

224
Q

What is oxidation?

A
  1. Loss of electrons.
  2. Increase in oxidation number.
225
Q

What is reduction?

A
  1. Electron gain.
  2. Decrease in oxidation number.
226
Q

What do all uncombined elements have as an oxidation number?

A

0

Examples include Zn, Cl₂, O₂.

227
Q

What do the oxidation numbers of the elements in a compound add up to?

228
Q

What is the oxidation number of a monoatomic ion equal to?

A

The ionic charge.

Eg if it’s Zn²⁺ then it is +2.

229
Q

What is the oxidation number of a polyatomic ion?

A

Adds up to the charge of the ion.

230
Q

State two differences between isotopes of the same element.

A
  1. Different number of neutrons.
  2. Different atomic masses.
231
Q

What is the order of the electron configuration?

A

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 4d¹⁰.

OCR A Chemistry (1 decks)

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