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Unit 3 - Elements and the Periodic Table

Elements, Compounds, and Mixtures

  • Changes
    • Physical Changes - Don’t produce a new substance
    • Chemical Changes - Produces a new substance
  • Matter Classification
    • Pure Substances
    • Can’t be broken into simpler compounds without going through a chemical change
      • Made of atoms that are chemically bonded to each other
    • Elements - Pure substances made of only 1 type of atom
    • Compounds - Pure substances made of 2+ types of atoms
    • Have fixed ratios between components
    • Mixtures
    • Mixing 2+ substances that are NOT chemically combined
    • Can be separated through physical means
      • Distillation - Separating components in a mixture through the use of their differing boiling points
      • Chromatography - Separating components using differences in their ability to pass through substrates
    • Don’t have fixed ratios between components
    • Types of Mixtures
      • Homogeneous Mixture - The components combined are indistinguishable
      • Heterogeneous Mixture - The components combined are distinguishable

Atomic Numbers and Electron Configurations

  • Quantum Orbitals
    • Orbitals - Location in an atom where an electron could be
    • An atom can have any number of orbitals depending on the number of electrons they have
    • Each orbital can hold 2 electrons
    • Quantum Number - Describes the location of an electron / describes the orbital
    • Three main quantum numbers used to describe orbitals: “N“, “L“, and “M“
      • N - Principal quantum number; describes the size of the orbital
      • Must be >0
      • You can think of this as what ring in the Bohr model the orbital coincides with
      • L - Angular momentum quantum number; describes the orbital shape
      • Can be spherical, dumbbell/peanut, clover, etc shaped
      • Can be between 0 → N-1
      • M - Magnetic quantum number; describes orientation of the orbital
      • Can be between -L → +L
    • Shells & Subshells
    • Electron shell - A group of orbitals with the same principle quantum number (N)
      • Shells are filled consecutively from the center/lowest energy orbitals outward
      • Different shells can hold different numbers of electrons
      • Full shells are the most stable
    • Electron subshells - A group of orbitals with the same principle quantum number (N) AND angular momentum quantum number (L)
      • Subshell Classifications
      • L = 0 → S Orbital
      • L = 1 → P Orbital
      • L = 2 → D Orbital
      • L = 3 → F Orbital
      • The number of different values the magnetic quantum number (M) can be is equal to the number of subshells of a certain classification
      • The number of orbitals is equal to the number of different combinations of N, L, and M (Can be calculated with N^2)
      • To calculate the number of electrons a shell can hold, you just double this number, since each orbital can hold 2 electrons
        • This can also be calculated with the formula 2N^2
    • Electron Configuration - How electrons are positioned in an atom
    • Orbital Notation - A diagram that shows shells, subshells, and orbitals using lines & arrows
      • Lines represent orbitals
      • Numbers & letters at the bottom represent the name of the orbital
      • Arrows represent electrons
      • Upward and downward arrows represent a M subscript s value of either 1/2 or -1/2
    • Pauli Exclusion Principle - No 2 electrons can have identical quantum numbers
      • A fourth quantum number, M subscript s represents the quantum spin of a number
      • Can have a value of either -1/2 and 1/2
      • Only 2 values → only 2 electrons can be in a orbital, otherwise at least 1 pair of electrons will have identical quantum numbers
    • Hund’s Rule - Electrons are placed in individual orbitals before being paired
    • Aufbau Principle - Electrons fill orbitals from lowest energy → highest energy
      • This means electrons fill from lower N to higher N
      • D and F are the exception; 3d has higher energy than 4s, so 4s will fill before 3d.
    • Follow the diagonal rule to determine order in which orbitals are filled

The History and Arrangement of the Periodic Table

  • Antoine Lavosier
    • Wrote “Elementary Treatise of Chemistry“ in 1789
    • Considered the world’s first modern chemistry textbook
    • Classified elements into 4 groups:
    • Acid-making
    • Gas-like
      • Wrongly classified light & heat as elements
    • Metallic
    • Earthy
      • Almost entirely made up of compounds
  • John Dobereiner
    • Arranged elements w/ similar properties into triads (groups of 3)
    • Difference between mass of elements 1 & 2 is about equal to difference in mass between 2 & 3
  • John Newlands
    • Arranged elements by atomic mass
    • Established “law of octaves“
    • Repeating pattern of similar properties every 8 elements
  • Dimitri Mendeleev
    • Created the first iteration of the modern periodic table
    • Arranged elements by atomic mass
    • Organized table rows/columns by chemical properties
  • Henry Moseley
    • Arranged elements by atomic number
    • Account for variation in natural isotopes
  • Periodic Table
    • Organized by atomic number (number of protons)
    • Columns have similar chemical properties due to having the same number of valence (outer) electrons
    • Each row is a new shell
    • Periods - A row on the periodic table
    • Atomic number increases from left to right
    • Chemical properties systematically change
    • Groups/Families - A column on the periodic table
    • All elements in groups have similar chemical properties
    • Cells - Give information about an element
    • Atomic number
    • Atomic mass
    • Atomic symbol
    • Element name
  • Elements
    • Natural: Elements 1-94
    • Man-Made: Elements 95-118
    • Metals: Left of the “staircase“ except hydrogen
    • Malleable
    • Ductile
    • Conduct heat & electricity
    • Mostly solids
    • Semi-Metals/Metalloids: The “staircase”
    • Properties of both groups
    • Non-Metals: Right of the “staircase“ plus hydrogen
    • Brittle
    • Poor Conductors
    • Can be any state
    • Main Group Elements
    • Alkali Metals - Group 1
      • Silver colored
      • Soft
      • Highly reactive with water/oxygen
      • Oxidizes in air
    • Alkaline Earth Metals - Group 2
      • Silver colored
      • More brittle than alkaline metals
      • Somewhat reactive
      • Low density, melting, and boiling points
    • Halogens - Group 17
      • Highly reactive w/ metals
      • Form salts
      • Toxic to most organisms
      • Mostly occur as diatomic molecules
    • Noble Gases - Group 18
      • Stable; don’t bond w/ other atoms
      • Non-flammable
      • Extremely low boiling points
      • Used in lights, produces colors when excited
    • Transition Metals
    • Form colored compounds
    • Some have unique properties
      • Some are magnetic
      • Some are very reactive
    • Inner Transition Metals
      • Can be radioactive
      • Lanthanides
      • Actinides

Electrons and the Periodic Table

  • Noble Gas Notation
    • Using noble gases to represent filled shells in longhand electron configuration
    • Separates valence and non-valence(core) electrons in an atom
  • Valence Electrons
    • The number of electrons on the outer shell of an atom
    • Determines the chemical properties of the atom
    • Correlated with the groups that the element is in in the periodic table
    • Group number = number of valence electrons
    • Determining Valence Electrons
    • Periods 1-3
      • Group number / highest S and P orbitals
    • Periods 4+
      • Highest S and P orbitals + partially filled d and f orbitals
  • Periodic Table & Orbitals
    • S-Block Elements - Elements in groups 1 & 2 + Helium
    • Has valence electrons in the S orbital
    • P-Block Elements - Elements in groups 13 → 18 - Helium
    • Rows 1-3
      • Has valence electrons in the S and P orbitals, with the last added electron being in the P orbital
    • Rows 4+
      • Has valence electrons in the S, D, and P orbitals, with “N“ of the D subshell being 1 less than the N of the S & P subshells and the last added electron being in the P orbital
    • D-Block Elements - Elements in groups 3 → 12 + Lutetium and Lawrencium
    • Has valence electrons in the S and D orbitals, with “N“ of the D subshell being 1 less than the N of the S subshell and the last added electron being in the D orbital
    • F-Block Elements - Lanthanides & Actinides - Lutetium and Lawrencium
    • Has valence electrons in the S and F orbitals, with “N“ of the F subshell being 2 less than the N of the S subshell and the last added electron being in the F orbital
    • Exceptions
    • Chromium
      • Predicted - [Ar] 4s^2 3d^4
      • Actual - [Ar] 4s^1 3d^5
    • Often D & F block elements that are transition metals
    • Happens because electrons fill lowest energy shell

Periodic Trends

  • Atomic Radius - 1/2 the distance between two identical atoms in a diatomic molecule
    • Increases down a group
    • Decreases across a row
    • More protons → electrons are pulled slightly closer together
  • Ionic Radius - Measure of the size of an ion
    • Anion - Negative ions (atoms that gain electrons)
    • Larger; More electrons cause more electron repulsion
    • Cation - Positive ions (atoms that lose electrons)
    • Smaller; less electrons cause less electron repulsion
    • Increase down a group
    • Decrease for cations across a row
    • Decrease for anions across a row
    • Increase when switching from cations to anions across a period
  • Ionization Energy - The energy required to remove an electron from an atom in a gas phase
    • Changes based
    • Nuclear charge
    • Distance from nucleus
    • The number of already removed electrons
    • First Ionization Energy - Energy needed to remove 1 electron from an atom
    • Second ionization energy - amount of energy to remove another electron after the first one is removed, etc
    • Main Group Elements
    • Increases across periods
    • Decreases down groups
    • Decreases between groups 2 & 13 and groups 15 & 16
  • Electron Affinity - Energy required to add an electron to a neutral atom in a gas phase
    • Decreases across a period
    • Increases down a group
  • Electronegativity - How much an atom attracts other electrons from other atoms
    • Increase across a period
    • Decrease down a group