Atomic Structure and The Periodic Table

States of Matter

  • Solid: Particles are closely packed, touching, and vibrate in fixed positions. Definite shape and volume.

  • Liquid: Particles are touching but vibrate more than solids and can move around. No fixed structure, adopts the shape of its container.

  • Gas: Particles are widely spaced and move randomly. They collide with each other and the container walls, bouncing off in all directions.

Changes of State

  • Melting: Solid to liquid; particles gain enough energy to break away from each other at a fixed temperature.

  • Freezing: Liquid to solid; particles lose energy, and forces holding them together become strong enough to solidify. Freezing point is the same temperature as melting point.

  • Boiling: Liquid to gas; particles gain enough energy to break away completely, forming bubbles in the liquid. Occurs at a fixed temperature dependent on atmospheric pressure.

  • Condensing: Gas to liquid; particles slow down due to cooling and stick together.

  • Evaporation: Liquid to gas; fast-moving particles on the surface break away.

  • Sublimation: Solid directly to gas upon heating; gas directly to solid upon cooling.

Atoms

  • Atoms and molecules explain chemical reactions.

  • The word "atom" comes from the Greek "atomos," meaning indivisible.

  • Democritus: Proposed that everything is made of tiny, indivisible particles called atoms, without physical evidence.

  • Plato & Aristotle: Believed the world was made of four elements: earth, water, fire, and air.

  • Robert Boyle & Isaac Newton: Revived the idea of atoms (corpuscles). Newton described them as tiny solid spheres, explaining gas diffusion but less successful in explaining chemical reactions.

  • John Dalton: Suggested different substances are made of different types of atoms with different weights. His assertion that atoms cannot be split was later proven incorrect.

  • Atoms are measured in nanometers (10-9 m) and range from 0.1-0.5 nm in diameter.

  • Scanning tunneling microscopy allows us to "see" atoms.

  • Manipulating atoms and molecules at this scale is the basis of nanoscience and nanotechnology.

Size of atoms and molecules

  • Grain of sand: ~ 0.1 cm

  • Human hair width: ~ 0.01 cm

  • Smoke particle: ~ 0.0001 cm

  • DNA width: ~ 0.0000001 cm

  • Atoms: ~ 0.00000001 cm

Determining Atom Size Experiment

  • Dissolve 1g of potassium manganate(VII) in 100 cm³ of water for a deep purple solution.

  • Dilute 10 cm³ of this solution to 100 cm³ with water, repeating until the color is barely visible.

  • Calculate the mass of potassium manganate(VII) in 1 cm³ of the most dilute solution where color is still visible.

  • Estimate the mass of potassium manganate(VII) in a drop of 1/1000 cm³.

Results, Calculations, and Formulas:

  • Solution: 1 cm³ of oleic acid in 1000 cm³ of petroleum ether.

  • Drop volume: 0.01 cm³.

  • Volume of oleic acid in 1 drop of solution: \frac{1}{1000} \\times 0.01 = 0.00001 cm³.

  • Thickness of oil patch = Volume/Area = volume/\pi r^2 = 0.00001 cm^3/(\pi \times r^2) cm2

Diffusion

  • Diffusion is the movement of a substance from high to low concentration.

Experiments Demonstrating Diffusion:

  • Potassium Manganate: A crystal is added to water, and the purple color spreads over time due to the movement of particles.

  • Perfume in Air: Measure the time it takes for the scent to travel different distances in the lab to estimate the speed of perfume molecules.

  • Bromine Vapor: Brown bromine vapor is placed below air; the gases mix due to diffusion, demonstrating the Kinetic Theory.

  • Ammonia and Hydrogen Chloride Gases: Gases react to form white ammonium chloride. \NH3(g) + HCl(g) \rightarrow NH4Cl(s)

Brownian Motion

  • Robert Brown observed random motion of pollen grains under a microscope.

  • Einstein explained Brownian motion using the Kinetic Theory.

Development of Atomic Structure

  • J.J. Thomson (1897): Plum pudding model – a sphere of positive charge with negative electrons scattered throughout.

  • Rutherford's Gold Foil Experiment:

    • Alpha particles fired at gold foil.

    • Expected: particles to pass straight through or slightly deflect.

    • Observed: some particles deflected at large angles or bounced back.

    • Conclusion: positive charge concentrated in a tiny nucleus.

  • Rutherford: Proposed most of the mass and positive charge is in a small nucleus, with electrons moving around it.

  • Niels Bohr (1913): Electrons move in stable orbits called shells or energy levels, using quantum physics.

  • James Chadwick (1932): Discovered the neutron, explaining extra mass in atoms.

Subatomic Particles

  • Proton:

    • Charge: +1

    • Mass: ~1 atomic mass unit (amu)

    • Location: Nucleus

  • Neutron:

    • Charge: 0

    • Mass: ~1 amu

    • Location: Nucleus

  • Electron:

    • Charge: -1

    • Mass: negligible

    • Location: Orbitals/shells around the nucleus

  • Atomic Number (Z): Number of protons, defines the element.

  • In neutral atoms: number of protons = number of electrons.

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

Isotopes

  • Isotopes are atoms of the same element with different numbers of neutrons.

  • Representation: ^A_Z X where A is the mass number, Z is the atomic number, and X is the element symbol.

Calculating Neutrons

  • Number of neutrons = Mass number (A) - Atomic number (Z)

Relative Atomic Mass (Ar)

  • Compares the mass of atoms to the ^{12}C isotope where \frac{1}{12} of ^{12}C = 1.

  • Definition: Weighted average mass of the isotopes of an element on a scale where carbon-12 has a mass of exactly 12 units.

Periodic Table

  • Arranged by increasing atomic mass, revealing patterns.

  • Mendeleev's table was better due to predicting undiscovered elements and their properties.

  • Modern table arranged by increasing atomic number.

  • Group: Vertical column of elements with similar properties.

  • Period: Horizontal row of elements.

Electron Arrangement

  • Determined by measuring the energy required to remove electrons.

Electronic Structure

  • Electrons are arranged in energy levels or shells around the nucleus.

  • Electrons fill the lowest available energy level first.

Rules for Electron Shells

  • 1st shell: max 2 electrons

  • 2nd shell: max 8 electrons

  • 3rd shell: max 8 electrons

  • 4th shell: 19th and 20th electrons go into this shell.

Dalton's Atomic Theory (1808) Revisited

  • All matter is composed of atoms: TRUE.

  • Atoms cannot be created, divided, or destroyed: FALSE (nuclear reactions).

  • All atoms of one element are alike: FALSE (isotopes).

  • Atoms of different elements have different weights: TRUE (but consider atomic mass).

  • "Compound Atoms" are formed by the combination of small whole numbers of atoms: TRUE (molecules).

Atomic Structure - Summary

  • Atoms have a dense nucleus containing protons (+ charge, ~1 amu) and neutrons (no charge, ~1 amu).

  • Electrons (- charge, negligible mass) move around the nucleus in shells.

  • The first shell holds up to 2 electrons, the second up to 8, and the third up to 8.

  • The charge of the nucleus is balanced by the electrons.

  • Aluminum (13 electrons): 2.8.3 electron configuration.

  • Atomic number (Z): number of protons, defines the element; also indicates number of electrons in a neutral atom.

  • Mass number (A): number of protons + neutrons.

  • Number of neutrons: A - Z.

  • Symbol: ^AZ X (e.g., chlorine with 18 neutrons: ^{35}{17}Cl).

  • Isotopes: atoms of the same element (same number of protons) with different numbers of neutrons.

    • Ar (relative atomic mass): average mass of an atom considering its isotopes and their proportions.

The Periodic Table

  • Elements arranged by atomic number (protons in the nucleus).

  • Elements with similar properties are in columns called groups.

  • Mendeleev predicted properties of undiscovered elements.

Key Contributors and Discoveries

Elements are created in stars. Elements up to iron (atomic number 26) are created through nuclear fusion in smaller stars. Elements with atomic numbers greater than 26 (iron) are created in supernovas. The four elements that the Ancient Greeks thought the world was made of were earth, water, fire, and air. Water is not considered to be an element because it can be broken down into simpler substances (hydrogen and oxygen). Metals known about 1600 AD included gold, silver, copper, iron, lead, tin, and mercury. Brandt obtained phosphorus from urine. It was special because it glowed in the dark (phosphorescence). Robert Boyle defined an element as a substance which cannot be broken into any simpler substances in 1661 AD. Humphry Davy used electrolysis at the start of the 19th century to isolate more elements. Davy isolated sodium, potassium, calcium, magnesium, strontium, and barium. Lithium, sodium, and potassium belong to the same chemical family (the alkali metals) because they have similar chemical properties and react in similar ways. At the start of the 19th century, the main problem that stopped people from being able to create a periodic table was the lack of knowledge about atomic weights and the properties of elements. Mendeleev listed the elements in order of increasing atomic weight. Mendeleev's table was more successful because he left gaps for undiscovered elements and predicted their properties. The modern periodic table is arranged in order of increasing atomic number. Gallium is used for in semiconductors, LEDs, and high-temperature thermometers.

The Modern Periodic Table

  • Groups: numbered 1 to 7 and then 0 (columns).

  • Periods: numbered 1-6 (rows).

  • Transition elements: middle section with no group number.

  • Group 1: Alkali metals.

  • Group 7: Halogens.

  • Group 0: Noble gases.

  • The position of the element reflects its underlying atomic structure.

Examples:

  • The element in group 4 and period 3 is Silicon (Si).

  • A group 4 metal is Titanium (Ti) or Zirconium(Zr).

  • The least reactive alkali metal is Francium (Fr).

  • A period 5 transition metal is Technetium (Tc).

  • The two elements which are liquid at room temperature are Bromine (Br) and Mercury (Hg)

  • The most reactive non-metal is Fluorine (F).

Electronic Configurations and Periodic Table

  • Lithium (Li) = 2.1

  • Boron (B) = 2.3

  • Oxygen (O) = 2.6

Q2. Electronic Configurations

  • Lithium: 2.1

  • Sodium: 2.8.1

  • Potassium: 2.8.8.1

  • Properties of Alkali Metals

*Aim: To learn about some of the properties of the alkali metals and how to use group trends to make predictions.

General Word Equation for a Group 1 Metal Reacting with Water:

  • Metal + Water → Metal Hydroxide + Hydrogen

Equations for the reactions of sodium and potassium with water:

  • Sodium: 2Na (s) + 2H2O (l) → 2NaOH(aq) + H2(g)

  • Potassium: 2K (s) + 2H2O (l) → 2KOH(aq) + H2(g)