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

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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
    • 

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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

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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

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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

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