Chemistry Ch 2

Group 8A - Noble Gases

• All gases at room temperature • Very low melting and boiling points • Very unreactive, practically inert • Very hard to remove electron from or give electron to

Nonmetals

• Found in all three states • Poor conductors of heat and electricity • Solids are brittle • Gain electrons and form anions • Upper right on the table - except H

Group 2A - Alkali Earth Metals

• Harder, higher melting point, and denser than alkali metals • Reactive, but less than alkali metals • Form stable, insoluble oxides

Group 1A - Alkali Metals

• Hydrogen though placed here - does not belong • Very reactive, never find uncombined in nature • Tend to form water-soluble compounds, e.g., NaCl, KCl

Group 7A - Halogens

• Nonmetals • F2 and Cl2 gases • Br2 liquid • I2 solid • All diatomic • Very reactive • React with metals to form ionic compounds, e.g., NaCl, NaBr

Metalloids

• Show mixed properties of both metals and nonmetals • Some are known as semiconductors because of their electrical conductivity

Metals

• Solids at room temperature, except Hg • Often have reflective surface - shiny • Conduct heat and electricity • Malleable - can be shaped • Ductile - can be pulled into wires • Lose electrons and form cations • About 75% of the elements are metals

Ernest Rutherford

discovered that there were three different kinds of emissions • alpha (α) rays • beta (β) rays • gamma (γ) rays

Two samples of carbon dioxide (CO2) are decomposed into their constituent elements. One sample produces 25.6 g of oxygen and 9.60 g of carbon, and the other produces 21.6 g of oxygen and 8.10 g of carbon. Show that these results are consistent with the law of definite proportions.

mass of oxygen/ mass of carbon 25.6g/9.60g= 2.61 mass of oxygen/ mass of carbon 21.6g/8.1= 2.61

Atomic number (Z) =

# protons

Mass Number (A) =

# protons + # neutrons

Atomic mass =

(fractional abundance of isotope)n x( mass of isotope)n

Main-group metals tend to:

- Lose electrons -Form cations with the same number of electrons as the nearest noble gas

Main-group nonmetals tend to:

-gain electrons -forming anions with the same number of electrons as the nearest noble gas

Dalton's Atomic Theory

1. Atoms can be neither created nor destroyed. 2. All atoms of an element are ALIKE in mass and other properties, but the atoms of one element differ from all other elements 3. Atoms of different elements combine in simple numerical ratio to form compounds 4. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way they are bound together.

Mole (mol) is the amount of material containing

6.02214 x 10(23) particles

Examples of Definite Proportions

A 100.0 g sample of sodium chloride contains 39.3 g of sodium and 60.7 g of chlorine mass of Cl / mass of Na = 60.7 g / 39.3 g = 1.54 A 200.0 g sample of sodium chloride contains 78.6 g of sodium and 121.4 g of chlorine mass of Cl / mass of Na = 121.4 g / 78.6 g = 1.54

Ions

Atoms can gain or lose electrons during chemical reactions

anions

Atoms that gain electrons form Example of gaining electrons: F + 1 e− → F−

cations

Atoms that lose electrons form Example of losing electrons: Li → Li+ + 1 e−

isotopes

Atoms with the same number of protons but different numbers of neutrons

Marie Currie and Henri Becquerel

Discovered that elements would emit small energetic particles and rays.

Millikan

Discovers oil drop

Thomson

Discovers the electron

Example of Multiple Proportions

Hydrogen and oxygen form the compounds water (H2O) and hydrogen peroxide (H2O2). The decomposition of a sample of water forms 0.125 g of hydrogen to every 1.00 g of oxygen. The decomposition of a sample of hydrogen peroxide forms 0.250 g of hydrogen to every 1.00 g of oxygen. Show that these results are consistent with the law of multiple proportions. mass oxygen to 1 g hydrogen peroxide/ mass oxygen to 1 g of water = 0.250 g/ 0.125g = 2

Law of Definite Proportions (law of constant composition)

If a compound is broken down into its constituent elements, the masses of the constituents will always have the same proportions regardless of their source or how they were prepared For example, the decomposition of 18.0 g of water results in 16.0 g of oxygen and 2.0 g of hydrogen, or an oxygen-to-hydrogen mass ratio of: mass ratio= 16.0 g o / 2.0 g H = 8.0

Law of Multiple Proportions

If two elements form more than one compound together, the masses of one element combined with a fixed mass of the second element form in ratios of small integers (always whole number) "Because there are twice as many oxygen atoms per carbon atom in carbon dioxide than in carbon monoxide, the oxygen mass ratio should be 2"

Law of Conservation of Mass

The total mass before a chemical reaction is the SAME as the total mass after the chemical reaction mass of reactants=mass of product