Atomic Structure and the Periodic table C1

Elements, Molecules Compounds & Mixtures definitions

• Element: A substance made of only one type of atom e.g He or O2

• Molecule: Two or more atoms covalently bonded together (can be from same or different element) e.g O2 or H2O

• Compound: Two or more different elements that are chemically bonded together in fixed proportions e.g H2O not O2

• Mixture: Two or more elements or compounds not chemically combined together e.g sea water (can be separated physically through filtration/distillation without a chemical reaction)

Subatomic Particles mass and charge

• Proton: Relative mass of 1, relative charge of +1, located in the nucleus.

• Neutron: Relative mass of 1, relative charge of 0 (neutral), located in the nucleus.

• Electron: Relative mass very small (0.0005), relative charge of -1, located orbiting the nucleus in energy levels (electron shells).

Atomic size and scale constants

• Atom radius: Approximately 1 x 10⁻¹⁰ metres (or 0.1 nanometres).

• Nucleus radius: Approximately 1 x 10⁻¹⁴ metres, which is 10,000 times smaller than the overall atom.

• Mass concentration: All the mass of an atom is concentrated inside its tiny nucleus, as it’s mostly empty space.

Atomic Number vs Mass Number rules

• Mass Number (Top): Total number of protons and neutrons in the nucleus.

• Atomic Number (Bottom): Number of protons in the nucleus (and electrons) , which determines what element they are (as if number changes it becomes a different element)

• In a neutral atom, the number of negative electrons always equals the number of positive protons (so they cancel out). To work out how many neutrons: Mass number - Atomic number

Ions formation and electron shells

• Ion is a charged particle formed when an atom gains or loses electrons to have a full outer shell.

• Metal atoms: Lose electrons from their outer shell to form positive ions (cations) with a full outer shell.

• Non-metal atoms: Gain electrons into their outer shell to form negative ions (anions) with a full outer shell.

Isotopes

• Isotopes: Different forms of the same element with the same number of protons (atomic number) but a different number of neutrons (mass number)- most elements exist naturally as a mixture of 2 or more isotopes with different masses e.g Carbon-12 and Carbon-13.

Relative atomic mass (Ar) and formula

The average mass of all the isotopes of an element, taking relative abundance into account (how common/ rare the isotope is in a %)

Relative atomic mass (Ar) = Sum of (Isotope abundance x Isotope mass) / Sum of abundance of all isotopes

Sum of abundances are equal to 100 when percentages

Electronic Structure 2,8,8 rule

• First shell: Holds a maximum of 2 electrons.

• Second shell: Holds a maximum of 8 electrons.

• Third shell: Holds a maximum of 8 electrons.

• Periodic Table link: The group number of an element tells you exactly how many electrons are in its outer shell, which gives the elements in that same group the same chemical properties (except Noble Gases unstable- as group 0 already have full outer shells).

History of the Atom timeline

• John Dalton (1803): Atoms are solid, spheres atoms that make up different elements.

• J.J Thomson (1897): Discovered electrons, creating the Plum Pudding Model that is a sphere of positive charge with negative electrons embedded (mass being evenly spread through whole atom and no empty space)

• Rutherford (1909): Nuclear Model with most of mass in a tiny nucleus containing positive charge surrounded by orbiting electrons with empty space.

• Niels Bohr (1913): Proved electrons orbit the nucleus at fixed energy levels (shells prevent atom from collapsing)

• Chadwick (1932): Discovered neutrons inside the nucleus.

Rutherford's alpha particle gold foil experiment

• Most alpha particles passed straight through the foil: Proves the atom is mostly empty space.

• Some alpha particles were deflected at small angle: Proves the nucleus of the atom has a concentrated positive charge that repelled them.

• A few alpha particles deflected straight back at a large angle: Proves the nucleus is tiny and contains most of the atom's mass.

Different mixtures examples that need separating

Mixture example: Insoluble solid that cannot dissolve in a solvent (water)

Mixture example: Soluble solid, liquid or gas (solute) that can dissolve in a solvent (water) forming a solution

Separating mixture technique: Filtration method

• Function: Separates an insoluble solid from a liquid solvent.

• Method: Place filter paper in a funnel over a flask. Pour the mixture in so the liquid passes through the pores as filtrate, while the solid particles are left behind as residue.

Separating mixture technique: Crystallisation/ Evaporation method

• Function: Separates a soluble solid from a liquid solvent (solution)

• Method: Heat the solution gently in an evaporating dish with water bath (more slow process), until it becomes saturated with some of the solvent evaporating (crystals start to form). Remove from heat and leave to cool slowly so large crystals grow, then filter and dry with filter paper so pure crystals (impurities filtered out with the remaining dissolved solution)

Evaporation is similar process: Heat solution in evaporating dish with a Bunsen burner, so all of the solvent evaporates and leaving dry crystals behind. Limitations: some of solids thermally decompose (breaking it down) and produced tiny crystals.

Separating mixture technique: Simple Distillation method

• Function: Separates a liquid solvent from a soluble dissolved solid solution, or two liquids with significantly different boiling points.

• Method: Heat the solution in a flask (with a bung so no gas escapes). The liquid with the lowest boiling point evaporates into vapour and pressure forces it into a water-cooled Liebig condenser, to condense back into a liquid, and is collected.

Limitations: only separates liquids with a large difference in boiling points (if they were to have the same b.p they would both evaporate and condense together) so fractional distillation is required.

Seperating mixture technique: Fractional Distillation method

• Function: Separates a mixture of multiple miscible liquids (mix into a single layer) with close boiling points (e.g. crude oil)

• Method: Heat the mixture (at a temperature of the lowest b.p) below a fractionating column packed with glass beads. Only the liquid with the lowest boiling point reaches the top first boiling into a vapour, so enters the condenser to condense back into liquid and collected. While if higher boiling point liquids evaporate early on once they reach the temperature cooler than their b.p they cool/ condense on beads and drip back down into flask.

Early Periodic Table development

• Early 1800s: Elements were arranged strictly in order of increasing atomic weight (now called atomic mass) because subatomic particles had not yet been discovered.

• Limitations : Placed completely dissimilar elements into the same columns because it ignored their chemical behaviours and incorrect as many elements not discovered yet.

Mendeleev's early Periodic Table improvements

• Kept elements arranged by atomic weight but swapped the order if their chemical properties did not fit the pattern.

• Left empty gaps in the table to keep elements with similar properties within same group and for undiscovered elements, allowing him to accurately predict the properties of the these missing elements.

Acceptance: Elements matching his exact predictions were later discovered and filled his gaps, which convinced other scientists his table was correct.

Modern Periodic Table arrangement

• Elements are arranged in order of increasing atomic (proton) number to resolve the anomalies in Mendeleev’s table.

Isotopes: Their discovery explained why ordering by atomic weight caused anomalies, as isotopes have different weights but the same chemical behaviours.

Has the group 0 and a seperate block for transition metals.

• Vertical columns are called Groups (elements have the same number of outer shell electrons and similar properties/ trends). Horizontal rows are called Periods (the same amount of electron shells).

Metals vs Non-Metals: General physical properties

• Metals: Good conductors of heat and electricity, malleable, shiny, sonorous, high melting and boiling points, and high density.

• Non-Metals: Poor conductors (insulators), brittle, dull appearance, low melting and boiling points, and lower density.

Metals vs Non-Metals: Chemical properties

• Metals: Found on the left and center of the periodic table, reacting to lose electrons and form positive ions (+).

• Non-Metals: Found on the top right side, reacting to gain or share electrons to form negative ions (-) or covalent bonds.

Group 1 Alkali Metals: Properties

• Features are different to general metals: Very soft to cut, low density (the first three float on water), and have low melting and boiling points that decrease down the group (as metallic bonding)

• Ion formed: Lose their single outer electron to form a +1 ion.

Group 1 Alkali Metals: Reactivity trend

• Trend: Reactivity increases going down the group.

• Explanation: As you go down, atoms have more electron shells (atoms radius increases), making the outer electron further away from the positive nucleus. The electron is more shielded from the positive charge of protons, so weaker electrostatic attraction requiring less energy, making it much easier to lose.

Group 1 Alkali Metals: Chemical reactions are vigorous

Alkali metal + Water→ Metal hydroxide + hydrogen gas

Alkali metal + Chlorine→ Metal alkali chloride (white solid salts)

Alkali metal + Oxygen→ Metal alkali oxide (tarnish instantly)

Group 7 Halogens: Properties

• Features: Non-metals that exist as diatomic molecules with covalent bonds (paired molecules like F₂, Cl₂, Br₂, I₂) and toxic so handled in fume cupboard.

• Trends: Melting and boiling points increase going down the group (as their states at room temp change from gas, liquids then solids)

• Ion formed: Gain one electron to form a -1 halide ion.

Group 7 Halogens: Reactivity trend

• Trend: Reactivity decreases going down the group.

• Explanation: As you go down, atoms have more electron shells (atoms radius increases), making the outer shell further away from the positive nucleus and more shielded. This reduces the electrostatic attraction from the positive charge of protons, making it harder to attract and gain an incoming electron.

Group 7 Halogens: Displacement Reactions

• Rule: A more reactive halogen displaces a less reactive halogen from an aqueous solution of its salt.

• Example: Chlorine + Sodium bromide → Sodium chloride + Bromine (Cl₂ + 2NaBr → 2NaCl + Br₂). Chlorine is more reactive than bromine, so it takes its place. These confirm that the reactivity decreases down the group.

Group 0 Noble Gases: Properties and trends

• Features: Colourless, non-flammable gases that exist as single monoatomic atoms.

• Reactivity: Completely unreactive (inert) because they have a stable full outer electron shells, meaning they have no tendency to lose or gain electrons.

• Trends: Boiling point and density increase going down the group due to an increase of electrons means stronger attractive force needing more energy to break.

Transition Metal: Properties

Typical metal properties e.g high boiling/ melting points

Can form more than one ion charges (Roman numerals represent these) they form different colours

Act as catalysts