Chapter 2 - Chem

Classical Laws: Law of Conservation of Mass (Chapter 2)

Mass/matter cannot be created or destroyed

• Mass at beginning and end of a chemical reaction must be the same

Classical Laws: Law of Definite Proportions (Constant Composition) (Chapter 2)

Compounds are composed of a definite (fixed) proportion by mass of their elements

• Use percent by mass of a compound to determine how much of each element a certain sample would contain

Classical Laws: Law of Multiple Proportions (Chapter 2)

If multiple compounds contain the same elements, the ratio of masses of them will be small whole numbers

• Can compare ratios of masses to determine chemical formulas

Definite Proportions: Percent Composition (Chapter 2)

If given masses of each element in a compound and mass of total compound, can find ratios or precents of each element in compounds

Ex.

A sample of CO₂ contains 47.13 g C and 125.57 g O. What is the percent by mass of each element?

• 47.13 + 125.57 = 172.70 g CO₂

• 47.13 g C/172.70 g CO₂ × 100% = 27.29% C

• 125.57 g O/172.70 g CO₂ × 100% = 72.71% O

Percent Composition Practice (Chapter 2)

A sample of CO₂ contains 47.13 g C and 125.57 g O. What is the percent by mass of each element?

• 47.13 + 125.57 = 172.70 g CO₂

• 47.13 g C/172.70 g CO₂ × 100% = 27.29% C

• 125.57 g O/172.70 g CO₂ × 100% = 72.71% O

Dalton’s Atomic Theory (Chapter 2)

Summarizes the classical laws:

• Matter is made of indivisible atoms (not quite true, they are made of subatomic particles)

• Each atom is identical to others of the same element but different from other elements (not quite true, isotopes and ions exist)

• Atoms combine in whole number ratios to form compounds

• Some atoms of the same elements can combine in different whole number ratios to form different compounds

History of the Atom (Chapter 2)

JJ. Thomson (1856-1940)

• Cathode ray = charged particle = electron (1897)

• Charge-to-mass ratio of electron (1897)

• “Plum pudding” model of atom (1904)

Robert Millikan (1868-1953)

• Charge and a mass of electrons (1909)

Henri Becquerel (1852-1908)

• Uranium emits rays that fog photographic film (1896)

Marie and Pierre Curie (1867-1934, 1854-1906)

• Radioactive elements

• Polonium and radium (1898)

Ernest Rutherford (1871-1937)

• a (alpha) and B (beta) particles (1998)

• Nuclear model of atom (1911)

• Proton (1920)

James Chadwick (1891-1974)

• Neutron (1932)

Subatomic Particles (Chapter 2)

Make up atoms

• Protons and Neutrons are inside the nucleus

• Electrons are outside of the nucleus but within the atom

Proton:

• Relative charge = +1

• Relative mass (amu) = 1

Neutron:

• Relative charge = 0

• Relative mass (amu) = 1

Electron:

• Relative charge = -1

• Relative mass (amu) = 0

Mass Number & Particles (Chapter 2)

Mass number: rounded atomic mass (decimal umber) from periodic table

• Sum of protons and neutrons

Atomic number: whole number from periodic table

• Number of protons (also electrons if neutral)

Example: sodium

• Mass number: 23

• Atomic number: 11

• Protons & electrons: 11

• Neutrons: 12 (23-11)

Mass Number Practice (Chapter 2)

Sodium

• Mass number:

• Atomic number:

• Protons & electrons:

• Neutrons:

• 23

• 11

• 11

• 12 (23-11)

Atomic Symbols (Chapter 2)

Mass number: 24 (12 protons +12 Neutrons)

Atomic number: 12 (12 protons)

Charge: 2+ (12 protons - 10 electrons)

Chemical element: Mg (Magnesium)

Particle & Atomic Symbols Practice (Chapter 2)

Tungsten mass number, atomic number, protons, neutrons, electrons

Atomic symbol for atom with 15 protons & 16 neutrons

How many protons, neutrons, & electrons in 73/32 Ge?

• 184 mass, 74 atomic

• 74 protons & electrons

• 184-74 = 110 neutrons

• 31/15 P

• p⁺ = 32, n⁰ = 41, e^- = 32

Ions (Chapter 2)

Ions are charge particles that result from an atom’s loss or gain of electrons

• Use number of protons and electrons to determine charge

• Remember: electrons are negative

Example:

A potassium ion contains 18 electrons, what is its charge?

• K has 19 protons if more protons than electrons (electron lost), +1

Ions Practice (Chapter 2)

An ion contains 16 protons, 16 neutrons, and 18 electrons. What element is it an ion of and what is its charge?

A magnesium ion contains 10 electrons, what is its charge?

• S is atomic number 16, if more electrons than proton (electrons gained), charge is negative, -2

• Mg has 12 protons, +2

Average Atomic Mass (Chapter)

Mass on periodic table is weighted average of the mass of all atoms of that element

Elements have isotopes- atom with different number of neutrons

Ex: C - 12 has 6 neutrons while C - 14 has 8 (number after dash = mass number)

Can use actual masses of isotopes (not quite whole numbers0 and their relative abundances (usually as precentages0 to calculate average atomic mass

The mass or precent of an isotope could also be calculated if given the average atomic mass

Calculate Average Atomic Mass Example (Chapter 2)

Percents are unknown, but together add up to 100%

Answer

Average Atomic Mass Practice (Chapter 2)

Answer

Average Atomic Mass Practice Pt.2 (Chapter 2)

Answer

Periodic Table Development (Chapter 2)

Mendeleev & Meyer independently developed ways to organize elements based on physical properties

Mendeleev gets most of the credit because he left gaps for where undiscovered elements should go - and they did, once discovered

Period Table Organization Pt.1 (Chapter 2)

Column: Group

Row: Period

Period Table Organization Pt.2 (Chapter 2)

Group numbers 1-18 for all, 1A-8A for representative elements, 1B - 8B for transition elements

Periodic Table Organization Pt.3 (Chapter 2)

Top left: Nonmetals

Metalloids: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Astatine

Right side/Middle: Metals

Row 4 - 7/Column 3 and so on: Transition Metals

Bottom section: Inner-Transition Metals

Main Group: 1, 2, 13, -18

Types of Elements (Chapter 2)

Metals

Metalloids

Nonmetals

Metals (Chapter 2)

Usually solid at room temperature, shiny, malleable, ductile, good conductors of heat & electricity, ex. copper, sodium, silver

Metalloids (Chapter 2)

Properties of both metals & nonmetals, semiconductors, ex. germanium, arsenic, boron

Nonmetals (Chapter 2)

May be gas, liquid, or solid at room temperature, brittle, poor conductors, ex. sulfur, nitrogen, bromine

Alkali Metals (Chapter 2)

Group 1 excluding (Hydrogen)

Very reactive

Very soft

Make water have a high pH Alkaline (Basic)

Alkaline Earth Metals (Chapter 2)

Group 2

Reactive

Soft

Make water alkaline (high pH Base)

Halogens (Chapter 2)

Group 17

Nonmetals & metalloids

Very reactive

2 gases, 1 liquid, 2 solids at room temperature

Not found in nature very easily

Noble Gases (Chapter 2)

Group 18

Nonmetals

Nonreactive (very stable)

Gases

Found in nature very easily