AS & A2 OCR Chemistry Notes

1.1.1 Atomic structures and isotopes

1. Isotopes
   
   • Isotopes are atoms of the same element with different numbers of neutrons and different masses
   • \(^{16}_{8}O\)
   • The upper left number is the mass number
   • The bottom left number is the atomic number
   • The letter is the chemical symbol

   Isotope	Protons,p+	Neutrons, n	Electrons, e-
   \(^{16}_{8}O\)	8	8	8

   • Isotopes can also be referred to by using simply the mass number or even just as element-mass.
2. Ions
   
   • An ion is a charged atom (positive or negative)
   • Positive ions are known as cations and have fewer electrons than protons \(\therefore\) they have a positive charge
   • Negative ions are known as anions and have more electrons than protons \(\therefore\) they have a negative charge

1.1.2 Relative mass

1. Carbon-12
   
   • Carbon-12 is the isotope around which relative masses of ions, atoms, and isotopes are defined \(\therefore\) the mass of a carbon-12 isotope is defined as exactly 12 atomic mass units

   Relative isotopic mass

   • Relative isotopic mass is the mass of an isotope relative to \(\frac{1}{12}\) th of the mass of an atom of carbon-12.

   Relative atomic mass

   • Relative atomic mass A_r is the weighted mean mass of an atom of an element relative to \(\frac{1}{12}\) th the mass of an atom of carbon-12.
   • Relative atomic mass can be found with an element in the periodic table along with its atomic number

   Determination of relative atomic mass

   • A mass spectromoter is used to find the percentage abundances of the isotopes in a sample of an element
   • You can then use the percentage abundances of the isotopes to calculate the relative atomic mass of an element with the following formula

   Relative atomic mass = \(\frac{(PercentageAbundance_{1} * IsotopicMass{1}) + (PercentageAbundance_{2} * IsotopicMass{2}) + ...}{100}\)

1.1.3 Polyatomic ions

1.1.4 Amount of substance and the mole

1.1.5 Chemical formulae

1. More relative masses
   
   • Some compounds exist as simple molecules whereas some exist as giant crystalline structures \(\therefore\) two terms are needed for relative mass, one for simple molecules and another for giant structures

   Relative molecular mass

   • The relative molecular mass M_r compares the mass of a molecule with the mass of an atom of carbon-12
   • M_r is calculated by summing the relative atomic masses of the elements which make up the molecule

   Relative formula mass

   • The relative formula mass compares the mass of a formula unit with the mass of an atom of carbon-12
   • Relative formula mass is calculated by summing the relative atomic masses of the elements in an empirical formula
2. Finding empirical formula from mass example
   

   In an experiment, 1.203g of calcium combines with 2.13g of chlorine to form a compound [A_r : Ca, 40.1; Cl, 35.5].

   • Step 1: Convert mass into moles using n = \(\frac{m}{M_{r}}\)

   n(Ca) = \(\frac{1.203}{40.1}\) = 0.030mol

   n(Cl) = \(\frac{2.13}{35.5}\) = 0.060mol

   • Step 2: To find the smallest whole-number ratio, divide by the smallest whole number

   n(Ca):n(Cl) = \(\frac{0.03}{0.030}\) : \(\frac{0.060}{0.030}\) = 1 : 2

   • Step 3: Write the empirical formula: CaCl₂

1.1.6 Moles and volumes

1. Worked example
   

   Calculate the mass of Na₂CO₃ required to prepare 100cm³ of a 0.250moldm^-3 standard solution

   • Step 1: Work out the amount in moles required

   n(Na2CO3) = \(\frac{c * V(cm^3)}{1000}\) = \(\frac{0.250 * 100}{1000}\) = 0.0250mol

   • Step 2: Work out the molar mass of Na₂CO₃

   n(Na2CO3) = \(23.0 * 2 + 12.0 + 16.0 * 3\) = 106.0gmol-1

   • Step 3: Rearrange n = \(\frac{n}{M_{r}}\) to calculate the mass of Na₂CO₃ required

   m = \(n * M\) = \(0.0250 * 106.0\) = 2.65g

   Converting between amount in moles and gas volumes

   n (mol) = \(\frac{volume (V)}{molar gas volume (V_{m})}\) \(\therefore\) at RTP (room temperature), n = \(\frac{V}{24}\)

   Ideal gas equation

   • p = pressure (Pa)
   • V = volume (m³)
   • n = amount of gas molecules (mol)
   • R = ideal gas constant = 8.31Jmol^-1K^-1
   • T = temperature (K)
   • pV = nRT

1.1.7 Reacting quantities

1.1.8 Acids, bases, and neutralisation

1.1.9 Acid-base titrations

1. Titration calculations
   
   • From the results of a titration, you will know both the concentration and reacting volume of one of the solutions
   • You will also know only the reacting volume of the other solution
   • Step 1: Work out the amount, in mol, of the solute in the solution for which you know both the concentration and the volume
   • Step 2: Use the equation to work out the amount, in mol, of the solute in the other solution
   • Step 3: Work out the unknown information about the solute in the other solution

1.1.10 Redox

1. Rules for elements
   
   • The oxidation number is always 0 for elements
2. Rules for compounds and ions
   
   • Each atom in a compound has an oxidation number
   • An oxidation number has a sign, which is placed before the number

   Combined element	Oxidation number	Examples
   O	-2	H2O, CaO
   H	+1	NH3, H2S
   F	-1	HF
   Na+, K+	+1	NaCl, K2O
   Mg2+, Ca2+	+2	MgCl2, CaO
   Cl-, Br- I-	-1	HCl, KBr, CaI2
   Special cases:
   H in metal hydrides	-1	NaH, CaH2
   O in peroxides	-1	H2O2
   O bonded to F	+2	F2O

   • Σ (Oxidation numbers) = total charge

   Roman numerals

   • Roman numerals are used in the names of compounds of elements that form ions with different charges
   • The Roman numeral shows the oxidation state (oxidation number) of the element, without a sign
   • e.g. iron(II) = Fe²⁺ with an oxidation number +2

   Redox reactions

3. Reduction and oxidation
   

   Reduction	Gain of electrons	Decrease in oxidation number
   Oxidation	Loss of electrons	Increase in oxidation number

1.1.11 Electron structure

1. s-orbitals
   
   • Sphere shape
   • Each shell contains one s-orbital
   • The greater the shell number \(n\), the greater the radius of its s-orbital
2. p-orbitals
   
   • Dumb-bnell shape
   • Three in each shell from \(n\) = 2 onwards at right-angles to one another
   • Referred to as p_x, p_y, p_z
3. d-orbitals and f-orbitals
   
   • Each shell from n = 3 contains five d-orbitals
   • Each shell from n = 4 contains seven f-orbitals

   Sub-shells

   • Within a shell, orbitals of the same type are grouped together as sub-shells

   Shell	s	p	d	f	Sub-shells present	Number of electrons in sub-shells	Number of electrons in shell
   1	1				1s	2	2
   2	1	3			2s + 2p	2 + 6	8
   3	1	3	5		3s + 3p + 3d	2 + 6 + 10	18
   4	1	3	5	7	4s + 4p + 4d + df	2 + 6 + 10 + 14	32

   Filling of orbitals

4. Orbitals fill in order of increasing energy
   
   • The sub-shells that make up shells have slightly different energy levels
   • Within each shell, the new type of sub-shell added has a higher energy
   • In the \(n\) = 2 shell, the order of filling is 2s, 2p
   • In the \(n\) = 3 shell, the order of filling is 3s, 3p, 3d
   • In the \(n\) = 4 shell, the order of filling is 4s, 4p, 4d, 4f
   • The 3d sub-shell is at a higher energy level than the 4s sub-shell \(\therefore\) the 4s sub-shell fills before the 3d sub-shell \(\therefore\) the order of filling is 3p, 4s, 3d

   Electrons pair with opposite spins

   • Electrons are negatively charged and repel one another
   • Electrons have a property called spin - either up or down
   • An arrow, \(\uparrow\) for up and \(\downarrow\) for down is used to indicate its spin
   • The two electrons in an orbital must have opposite spin to help counteract the repulsion between the charges of the two electrons

   Orbitals of the same energy are occupied singly first

   Electron configuration

   • Electron configuration can be written as 1s^a2s^b2p^c… etc.
   • e.g. For krypton (Z = 36) the electron configuration would be 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶
   • Electron configuration can also be expression in terms of the previous noble gas in the periodic table
   • e.g. For Li the electron configuration could be written as 1s²2s{1} or more simply [He]2s¹
   • Cations are formed when atoms lose electrons
   • Anions are formed when atoms gain electrons

2.2.1 How Fast

1. Arrhenius equation
   
   • k = Ae^-E_a/RT

2.2.2 Acids and Bases

1. Finding pH of a strong base
   
   1. Calculate [OH^-]
   2. Use K_w to find [H⁺]
   3. Calculate pH
2. Buffer solutions and weak acids
   
   1. pH of a weak acid
      
      • e.g CH₃COOH \(\rightleftharpoons\) H⁺ + CH₃COO^-
      • HA \(\rightleftharpoons\) H⁺ + A^-
   2. pH of a buffer solution
      
      • HA + NaOH \(\rightarrow\) NaA + H₂O

      	HA	NaOH	NaA
      Start	0.5	0.2	0
      Change	-0.2	-0.2	+0.2
      End	0.3	0	0.2

   3. Artificial construction of a buffer solution
      
      1. HA \(\rightarrow\) vol 25cm³ and [0.1] = \(\frac{25 * 0.1}{1000}\)
      2. A^- \(\rightarrow\) vol 10cm³ and [0.1] = \(\frac{10 * 0.1}{1000}\)
      3. New conc. [HA] = \(\frac{2.5 * 10^{-3} * 1000}{35}\) = 0.0714
      4. [A^-] = \(\frac{1*10^{-3} * 1000}{35}\) = 0.0236
      5. K_A = \(1*10^{-4}\)
      6. \(\frac{K_{A} * [HA]}{[A^{-}]}\) = 3.6
3. Water
   
   • H₂O \(\rightleftharpoons\) H⁺ + OH^- ; Δ H = +ve
   • While the pH of water changes with temperature, water cannot become more acidic or alkaline. Instead, the position of neutral changes.

2.2.3 Entropy