Year 11 Chemistry Ultimate Summary By Shereef Magar

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How to Use This Guide

1. Use this as a quick recap, not a full textbook replacement
2. Focus first on definitions, formulas, and key trends
3. DO A PAST PAPERS!! click here to download some
4. If you want a more indepth sheet go to mine over here

Atom

• Definition: An atom is the smallest unit of an element that makes up all matter.
• Subatomic Particles: Inside an atom, there are three subatomic particles: protons (+), neutrons (0), and electrons (-).
• Nucleus: The dense positively charged centre of an atom that contains protons and neutrons. (Plural is nuclei).
• Electron Cloud: A large empty space around the nucleus where the electrons roam around.
• Mass: Protons and neutrons have the same mass; electrons have a drastically lower mass.

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Nuclear Notation / Isotopic Notation

${}_{Z}^{A}X$

Symbol	Definition
X	Symbol of element
A	Mass Number: Total number of neutrons and protons
Z	Atomic Number: Number of protons

Notes: Charge indicates if there are more or less electrons. Example: ${}_{8}^{18}O^{2-}$.

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Ions

• Definition: An ion is an atom (or group of atoms) that has gained or lost electrons so that it has a charge.
• Cation: A positively charged ion. It forms when an atom loses one or more electrons.
• Anion: A negatively charged ion. It forms when an atom gains one or more electrons.

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Electrostatic Force

• Electrostatic Attraction: The attraction between oppositely charged particles.
• Electrostatic Repulsion: The repulsion between same charged particles.

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Isotopes

• Definition: Atoms of the same element with the same protons but different neutrons (and therefore different mass number).
• Chemical Properties: Isotopes have the same chemical properties because they have the same amount of protons and electrons (same element).
• Physical Properties: Isotopes have different physical properties because they have a different number of neutrons and therefore different masses.
• Impact of Mass: Difference of mass can affect density, melting/boiling points, and stability.
• Instability: If an atom has relatively too many neutrons compared to protons, the atom can become unstable and can release radiation in order to become more stable.

Calculation: Number of Neutrons

• $\text{Number of Neutrons} = \text{Atomic Mass} - \text{Atomic Number}$.

Isotopes of Carbon

• Carbon-12 ($98.9\%$): ${}_{6}^{12}C$. Neutrons $= 12 - 6 = 6$.
• Carbon-13 ($1.1\%$): ${}_{6}^{13}C$. Neutrons $= 13 - 6 = 7$.
• Carbon-14 ($0.0001\%$): ${}_{6}^{14}C$. Neutrons $= 14 - 6 = 8$.

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Atomic Models & History

John Dalton

• All matter is made up of atoms.
• All atoms of one element are identical (not correct because of isotopes).
• Atoms of different elements are different to each other.
• Atoms combine in whole number ratios to form compounds.
• Atoms cannot be divided (not correct because atoms can be divided into subatomic particles).

J.J. Thomson: Plum Pudding Model

• Showed that the atom was made up of a positive nucleus with negatively charged electrons spread throughout like plums in a pudding.
• Discovered the Electron.

Rutherford Experiment

• Method: Alpha particles (positively charged) were fired at a thin sheet of gold foil.
• Observation: Majority of particles passed through, but a few were deflected back.
• Conclusions:
  1. Majority of the atom is empty space.
  2. There is a positive central region containing most of the mass, named the nucleus. (Like charges repel).

Bohr Model

• Electrons orbit the nucleus in fixed energy levels (shells).
• Shell Energy: The closer the shell is to the nucleus, the lower its energy (radius is smaller). The further away, the higher its energy (radius is larger).
• Movement: Electrons occupy one shell at a time and cannot exist between levels.
• Excited State: If an electron gains energy (absorbs a photon), it jumps to a higher level.
• Ground State: If an electron loses energy (releases a photon), it falls back to a lower level.
• The energy of the photon is equal to the difference in energy between the two levels.

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Mass Spectrometry

1. Vaporization: Element is heated until it becomes a gas. (Often assumed pre-done).
2. Ionisation: Vaporized element passes through an electron beam, losing electrons to form $+1$ ions (rarely $+2$).
3. Acceleration: Ions are accelerated by an electric field toward a negative plate.
4. Deflection: Ions pass through a magnetic field and deflect based on their mass to charge ratio ($m/z$).
   • Heavier isotopes deflect less; lighter isotopes deflect more.
   • Only charged particles can be deflected.
5. Detection: The detector measures the number of ions reaching it at each $m/z$ value to determine relative abundance and mass.

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Relative Atomic and Isotopic Mass

• Relative Atomic Mass ($A_r$): The weighted average mass of all isotopes of an element compared to $1/12th$ of the mass of a carbon-12 atom. Also called Atomic Weight.
• Relative Isotopic Mass: The mass of an individual isotope compared to $1/12th$ of carbon-12.
• Relative Abundance: The percentage abundance of an isotope.

Calculations

1. Finding $A_r$ from Abundance $A_r = ( \text{mass}_1 \times \text{abundance}_1 ) + ( \text{mass}_2 \times \text{abundance}_2 )$

• Example Chlorine: $(34.969 \times 0.75) + (36.966 \times 0.25) = 35.5$.

2. Finding Abundance from $A_r$ $A_r = \frac{(m_1 \times p_1) + (m_2 \times (100 - p_1))}{100}$

• Example Carbon: $A_r = 12.20$, Isotopes Carbon-12 and Carbon-13.
  • $12.20 = \frac{(12 \times p_1) + (13 \times (100 - p_1))}{100}$.
  • Result: $P1=80$, $P2=20$. Carbon-12 is $80\%$, Carbon-13 is $20\%$.

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Periodic Table Trends

• Atomic Radius: Distance from nucleus to valence electrons.
• First Ionisation Energy: Energy required to remove one electron.
• Electronegativity: Ability of an atom to attract shared electrons (Noble gases excluded).

Trend	Down a Group (Decreased Attraction)	Across a Period (Increased Attraction)
Reasoning	Increased shells and shielding	Increased core charge (protons)
Atomic Radius	Increases	Decreases
Ionisation Energy	Decreases (easier to remove)	Increases (harder to remove)
Electronegativity	Decreases	Increases

Metallic and Non-Metallic Character

• Metallic Character: Ability to lose electrons to form a positive ion. Metals have low ionisation energy and low electronegativity.
• Non-metallic Character: Ability to gain electrons to form a negative ion. Non-metals have high ionisation energy and high electronegativity.

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Bonding

Type	Structure	Bonding Mechanism
Metallic	3D lattice of positive ions in a “sea of delocalized electrons”.	Attraction between positive metal ions and delocalized valence electrons.
Ionic	3D rigid lattice of alternating positive and negative ions.	Attraction between positive metal ions and negative non-metal ions.
Covalent Molecular	Molecules with strong intramolecular and weak intermolecular bonds.	Attraction between positive nuclei of non-metals and shared electrons.
Covalent Network	Giant 3D continuous lattice of non-metal atoms.	Continuous strong covalent bonds throughout the structure.

Physical Properties

• Melting/Boiling Points:
  • High: Metallic, Ionic, Covalent Network (requires large energy to break strong bonds).
  • Low: Covalent Molecular (weak intermolecular bonds easily broken).
• Electrical Conductivity: Requires free-moving charged particles (electrons or ions).
  • Metals: Conduct as solid/liquid due to delocalized electrons.
  • Ionic: Conduct only when molten or aqueous (ions are free); not as a solid.
  • Covalent: Usually non-conductive (except graphite).
• Malleability/Ductility (Metals): Metal ions slide over each other without breaking the non-directional bond with the electron sea.
• Brittleness:
  • Ionic: Force causes like-charges to align and repel, shattering the lattice.
  • Covalent: Directional bonds break and cannot easily reform.

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Carbon Allotropes

Allotropes are different structural forms of the same element with different properties.

• Diamond: Each Carbon bonded to four others (tetrahedral). Extremely hard; high MP/BP; non-conductive.
• Graphite: Each Carbon bonded to three others in hexagonal layers. Conductive (one delocalized electron per C). Soft/slippery (layers held by weak forces).
• Graphene: A single layer of graphite.
• Fullerenes: Hollow cage structures (nanomaterials).
  • Buckyball ($C_{60}$): Soccer ball structure. Low MP/BP; poor conductor.
  • Carbon Nanotubes: Cylindrical. High tensile strength; high MP/BP; good conductor.

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Nanomaterials

• Scale: At least one dimension between $1-100~nm$. Larger materials are “bulk”.
• Surface Area: Very large surface-area-to-volume ratio increases reaction rates (catalysts).
• Properties: Can differ from bulk materials in colour, MP, and reactivity.
• Safety: Small size allows absorption into cells, potentially causing respiratory issues.

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Spectra and AAS

• Absorption Spectrum: White light passes through a substance; electrons absorb specific photons and move to excited states. Displays as black lines on a continuous spectrum.
• Emission Spectrum: Excited electrons release photons to return to ground state. Displays as coloured lines on a black background. Unique to each element.
• Flame Test: Similar to emission; photons change the flame’s colour.
• Atomic Absorption Spectroscopy (AAS): Used to find metal concentration in solution. Light from a hollow cathode lamp is absorbed by free metal atoms in a flame; more absorption = higher concentration.

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Energy in Reactions

• Exothermic: Releases energy ($\Delta H$ is negative). Bonds forming > Bonds breaking. Temperature increases.
  • Examples: Combustion, Neutralisation.
• Endothermic: Absorbs energy ($\Delta H$ is positive). Bonds breaking > Bonds forming. Temperature decreases.
  • Examples: Photosynthesis, Ionisation.
• Activation Energy ($E_a$): Minimum energy required for a reaction to start.
• Enthalpy ($H$): Total energy in a system (kinetic + potential).

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Mixtures and Separation

• Pure Substance: One type of particle (Element or Compound).
• Mixture: Physically combined; components retain properties.
  • Homogeneous: Evenly distributed (e.g., aqueous solutions).
  • Heterogeneous: Not evenly distributed.

Techniques

• Sieving/Filtration: Based on particle size.
• Decantation/Separating Funnel: Based on density.
• Evaporation/Distillation: Based on boiling points.
• Fractional Distillation: Separates complex mixtures (like crude oil) using a column; lower BP condense higher up.

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Organic Chemistry

• Alkanes: Single bonds only ($C_nH_{2n+2}$). Saturated.
• Alkenes: At least one double bond ($C_nH_{2n}$). Unsaturated.
• Benzene: $C_6H_6$ ring with delocalized electrons.

Reactions

• Addition (Alkenes): Double bond breaks to add new atoms. Orange bromine turns colourless.
• Substitution (Alkanes/Benzene): One atom replaced by another. Requires UV light or catalysts.
• Combustion: Complete ($+ O_2 \to CO_2 + H_2O$) or Incomplete ($\to CO + H_2O$).

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Stoichiometry

• Mole Formula: $n = \frac{m}{M}$.
• Particles: $N = n \times 6.022 \times 10^{23}$.
• Percent Composition: $\frac{\text{Molar Mass of element}}{\text{Molar Mass of Compound}} \times 100$.
• Percentage Purity: $\frac{\text{Mass(pure)}}{\text{Mass(impure)}} \times 100$.

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Scientific Enquiry

• Primary Data: Collected yourself.
• Secondary Data: Collected by others (textbooks, etc.).
• Variables: Independent (X-axis, what you change); Dependent (Y-axis, what you measure).
• Reliability: Increased by more trials to detect outliers and reduce random error.

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Successive Ionisation Energy

• Increases as each electron is removed from an increasingly positive ion.
• Significant jumps occur when removing from a new inner shell (closer to nucleus).

Species Data Example

Species	Symbol	Protons	Neutrons	Configuration
V	${}^{24}Mg^{2+}$	12	12	2, 8
W	${}^{1}H$	1	0	1
X	${}^{33}S^{2-}$	16	17	2, 8, 8
Y	${}^{40}Ar$	18	22	2, 8, 8
Z	${}^{31}P$	15	16	2, 8, 5

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