Science 1st quarter

Matter: solid, liquid, gas, plasma, and Bose-Einstein condensate.

Matter - occupies space and have mass, filled with atoms.

Atoms - the smallest and simplest particle. Composed of:

• Protons - positive charge

• Neutron - no charge

• Electron - negative charge

• Atomic Number = Number of Protons = Number of Electrons

• Atomic Mass = Number of Proton + Number of Neutron

• Neutron = Atomic Mass - Atomic Number

Atomic Theory Timeline

• Democritus - a Greek philosopher who introduced “Atom”. The term "atom" comes from the Greek word "atomos" (ἄτομος), which means "uncuttable" or "indivisible."

• John Dalton – Said atoms are tiny, solid balls; first modern atomic theory.

• J.J. Thomson – Discovered electrons; made the “plum pudding” model.

• Ernest Rutherford – Found the nucleus; atoms are mostly empty space.

• Niels Bohr – Said electrons move in fixed orbits around the nucleus.

• Erwin Schrödinger – Said electrons move in clouds, not fixed paths.

Periodic Table of Elements

Periodic Table of Elements - Arrangement of these elements by increasing atomic number and similar properties.
Elements (Atoms) - are substances that cannot be broken down any further by chemical means.

Dmittri Mendelev - is a Russian chemist who is the father of the modern periodic table. He originally arranged the elements according to increasing mass.

Henry Moseley - corrected Dmitri Mendeleev's Periodic Table by arranging elements by atomic number (number of protons), not atomic mass.

1. Periods (→ Rows, left to right)

   • Go horizontally across the table

   • Tell the number of energy levels (shells) an atom has

   • There are 7 periods

2. Groups or Families (↓ Columns, top to bottom)

   • Go vertically down the table

   • Elements in the same group have the same number of valence electrons

   • Same chemical properties

   • There are 18 groups

Types of the Elements in Periodic Table of Elements

• Metals - Shiny, can be bent, good conductors of heat and electricity

• Alkaline Earth Metals (Group 2) - Reactive but less than Group 1, form basic solutions in water

• Transition Metals (Middle block) - Hard, shiny, good conductors and can form colored compounds

• Metalloids - Have properties of both metals and nonmetals

• Nonmetals - Dull, brittle, don’t conduct heat/electricity well

• Halogens (Group 17) - Very reactive nonmetals and often form salts with metals

• Noble Gases (Group 18) - Stable, unreactive, used in lights

Atomic Properties

• Ionization Energy – how much energy is needed to remove an electron

• Electron Affinity – how much energy is released when an atom gains an electron

Electron Configuration

Electron Configuration - the distribution of atoms in an atoms’ energy levels, and describes how electrons are arranged around the nucleus

Energy Levels and Sublevels

• Energy Levels - numbers 1, 2, 3, etc.

• Sublevels - can hold a specific number of electrons and has s, p, d, f within each energy level (Sharp, Principal, Diffuse, Fundamental)

  • S = 2e

  • P = 6e

  • D = 10e

  • F = 14e

1. Aufbau Principle - Electrons fill the lowest energy levels first. Start filling from 1s
   ightarrow 2s
   ightarrow 2p
   ightarrow 3s… (lowest to highest)

2. Pauli Exclusion Principle - Only 2 electrons per orbital, and they must have opposite spins.

3. Hund’s Rule - Electrons fill orbitals one at a time before pairing. In the same sublevel (like 3p), one electron goes into each orbital first.

Period = Row

Family = add the two last exponents

Lewis Electron Dot Structure (LEDS)

Named after Gilbert N. Lewis (1875-1946)

• Consists of an element symbol with dots placed around it

• The symbol represents the kernel (nucleus + inner electrons)

• The dots represent valence electrons (electrons in the outermost shell)

• Each dot = 1 valence electron

• Shows the valence electrons of an atom

• Helps in understanding chemical bonding between atoms

Reminders for drawing Lewis Dot Structures

• One dot per side, either clockwise or counterclockwise

• Place dots one at a time on each side before pairing

• You only pair the first side

• Maximum of 8 dots

Valence electrons - the electrons found in the outermost shell (energy level) of an atom.

• Determine chemical reactivity of an element

• Involved in bonding with other atoms (ionic or covalent)

• The number of valence electrons tells us how many bonds an atom can form

Octet Rule - states that atoms tend to gain, lose, or share electrons to get 8 valence electrons, making them stable (like noble gases).

• Explains why atoms bond

• Helps predict how atoms will bond and what kind of molecules they form

• Leads to stable electron arrangements

Chemical Bonding

Chemical Bonding - the force that holds atoms together
Valence Electron - electrons in the highest occupied energy level

Different ways to bond

• Gaining Electron

• Losing Electron

• Pooling Electron

• Sharing Electron

Types of Chemical Bonding

Ionic Bond

• Ionic Bond - involves the transfer of electron from one atom

• Bond between metal and nonmetal

• Forms ionic compound

• Electronegativity difference is greater than 2.0

• Ion - atom that is no longer neutral since it has lost or gained an electron

  • Cation - lost electrons, positive charge

  • Anion - gained electrons, negative charge

Naming Ionic Compounds

• Binary Compound - contain only 2 elements

  • Rules:

    1. Write the name of the metal ion

    2. Write the name of the nonmetal ion (change the ending to -ide)

• Binary Compound Contain Transition Group B (transition) - have more than one possible oxidation state.

  • Exception: Silver, Zinc, Cadmium

  • Two methods that can clarify when more than one charge is possible:

    • Stock System - roman numeral

    • Classical Method - suffixes

• Ternary Compound - contain 3 different elements, often involve polyatomic ions

• Polyatomic Ion - group of atoms that are bonded together with a charge. Often end with -ate or -ite

Covalent Bond

• Covalent Bond - pairs of e- are shared between nonmetal atoms

• Electronegativity difference 2.0

• Partake polyatomic atoms

Types of Covalent Bond

• Single Covalent Bond

• Double Covalent Bond

• Triple Covalent Bond

• Polar Covalent Bond

  • Polar - affects how chemicals mix

• Non Polar Covalent Bond

• Coordinate Covalent Bond

Naming Covalent Compounds

• Typically composed of two nonmetals

• Rule:

  1. Name the first element

  2. Name the second element with the suffix “-ide”

  3. Use prefixes to denote the number of each type of atom (Mono is never used to name the first element)

Acid Naming

Acid - starts with hydrogen and ends with a nonmetal or polyatomic ion.

• Binary Acid - made from hydrogen and a single element

  • “Hydro” + element name + “ic” Acid

• Oxyacid - made from hydrogen and oxygen containing polyatomic ion

  • Polyatomic Ion Name + “Acid”

  • Change “ate” ending to “ic”

  • Change “ite” ending to “ous”

Metallic Bond

Metallic Bond - formed when positively charged kernels are attracted to delocalized electrons

• Metal is the only substance involved

Percentage Composition

Percentage composition – shows the percent by mass of each element in a compound; used to understand how much each element contributes to the total.

Formula to use:
ext{Percentage Composition} = rac{ ext{Mass of element in 1 mole of compound}}{ ext{Molar mass of compound}} imes 100

Steps:

1. Find molar mass of compound.

2. Find the total mass of each element.

3. Plug into the formula and multiply by 100.

Example – H_2O (Water):

• H: 2 imes 1.01 = 2.02 ext{ g}

• O: 1 imes 16.00 = 16.00 ext{ g}

• Molar mass = 18.02 ext{ g}

• Hydrogen = (2.02 ext{ g} ext{ H} ext{atoms} ext{ / } 18.02 ext{ g} ext{ H}_2 ext{O}) imes 100 = 11.21 ext{\%}

• Oxygen = (16.00 ext{ g} ext{ O} ext{atoms} ext{ / } 18.02 ext{ g} ext{ H}_2 ext{O}) imes 100 = 88.79 ext{\%}

Empirical Formula

Empirical formula – simplest whole‑number ratio of atoms in a compound.

Molecular Formula

Molecular formula – actual number of atoms in a molecule.

To find empirical formula:

1. Assume 100 ext{ g} → percent = grams.

2. Convert grams to moles.

3. Divide all mole values by the smallest mole value.

4. Multiply to get whole numbers if needed.

**Example: **40 ext{\%} **C, **6.7 ext{\%} **H, **53.3 ext{\%} O

• C: 40 ext{ g} ext{ C} ext{ / } 12.01 ext{ g/mol} = 3.33 ext{ mol}

• H: 6.7 ext{ g} ext{ H} ext{ / } 1.01 ext{ g/mol} = 6.63 ext{ mol}

• O: 53.3 ext{ g} ext{ O} ext{ / } 16.00 ext{ g/mol} = 3.33 ext{ mol}

• Divide all by 3.33 → C=1, H=2, O=1

• Empirical formula = CH_2O

To find molecular formula:

n = rac{ ext{Molar mass (given)}}{ ext{Molar mass of empirical formula}}

• Empirical molar mass (CH_2O) = 30 ext{ g/mol}

• If compound’s molar mass = 180 ext{ g/mol} → $$n =