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