Chemistry Unit 1

Mercury (I)

Hg₂²⁺

Ammonium

NH₄⁺

Nitrite

NO₂^-

Nitrate

NO₃^-

Sulfite

SO₃^2-

Sulfate

SO₄^2-

Hydrogen Sulfate/Bisulfate

HSO₄^-

Hydroxide

OH^-

Cyanide

CN^-

Phosphate

PO₄^3-

Hydrogen Phosphate

HPO₄^2-

Dihydrogen Phosphate

H₂PO₄^-

Thiocyanate

NCS^-/SCN^-

Carbonate

CO₃^2-

Hydrogen Carbonate/Bicarbonate

HCO₃^-

Hypochlorite

ClO^-/OCl^-

Chlorite

ClO₂^-

Chlorate

ClO₃^-

Perchlorate

ClO₄^-

Acetate

C₂H₃O₂^-

Permanganate

MnO₄^-

Dichromate

Cr₂O₇^2-

Chromate

CrO₄^2-

Peroxide

O₂^2-

Oxalate

C₂O₄^2-

Thiosulfate

S₂O₃^2-

Applications of chem before 1000 b.c

• Use of embalming fluids

• Processing of natural ores to produce metals for ornaments and weapons

Greek view of Chem (400 bc)

• Proposed that matter was composed of four fundamental substances (earth, fire, air, and water)

• Considered the question of whether matter is infinitely divisible or is composed of small, indivisible particles

Alchemists contributions to chem

• Helped discover several elements 

• Learned to prepare mineral acids

Georg Bauer

developed systematic metallurgy (16th century)

Paracelsus

• discovered the medicinal application of minerals 

Robert Boyle

• quantitative experiments measure pressure and volume of air

• first definition of element

  • A substance = element unless it can be broken down into two or more substances

• Metals not elements, thought method to change one metal to another will be found

Joseph priestly

discovered oxygen gas

Georg Stahl

• Suggested phlogiston came from burning material 

• thought substances that burn in closed containers stop burning since the air in the container is saturated with phlogiston

• Oxygen = low in phlogiston bc supports combustion

  • Was originally called dephlogisticated air

Antoine Lavoisier

• Explained combustion

• measurement = essential operation of chemistry

• Verified the law of conservation of mass 

• Conducted experiments showed that combustion involved oxygen

• Discovered that life supported by a process involving oxygen, similar to combustion

Law of Conservation of Mass

• Mass is neither created nor destroyed in a chemical reaction 

Joseph Proust

• Proposed the principle of the constant composition of compounds or Proust’s law or the law of definite proportion

law of definite proportion

• A given compound always contains exactly the same proportion of elements by mass

John Dalton

• proposed law of multiple proportions; dalton’s atomic theory

• prepared first table of atomic masses (many proved to be wrong)

law of multiple proportions

When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 g of the first element can always be reduced to small whole numbers

Dalton’s Atomic theory

• elements made up of atoms

• Atoms of a given element are identical

• Chemical compounds when atoms of different elements combine with each other

  • compound always has same relative numbers and types of atoms

• Chemical reactions = reorganization of the atoms

  • Atoms are not changed

Atomic masses

total mass of Atom

Joseph Gay-Lussac

Measured the volumes of gases that reacted with each other under the same temperature and pressure conditions

Avogadro’s hypothesis

• At the same temperature and pressure, equal volumes of different gases contain the same number of particles

  • Makes sense if the distances between the particles in a gas are very great compared with the sizes of the particles

  • Volume of a gas is determined by the number of molecules present, not by the size of the individual particles

What is not discussed in dalton’s atomic theory

Isotopes, protons, neutrons, electrons, ions

J.J Thomson

• Studied electric discharges in cathode-ray tubes

• cathode ray = negatively charged particles (electrons)

  • Cathode ray was produced at the negative electrode when high voltage was applied to the tube 

    • Repelled by the negative pole of an applied electric field

• Determined the charge-to-mass ratio of an electron

• e - Charge on the electron (in coulombs)

• m - Electron mass (in grams)

JJ thomson assumptions

• All atoms must contain electrons

  • Electrons can be produced from electrodes made of various metals

• Atoms must contain some amount of positive charge 

  • Atoms were known to be electrically neutral

Plum Pudding Model

Atoms consist of a diffuse cloud of positive charge with the negative electrons embedded randomly in it

Robert Millikan

• Performed experiments involving charged oil drops, which helped determine the magnitude of electron charge 

  • Used this value and the charge-to-mass ratio to calculate the mass of an electron as 9.11 ×10^–31 kg

Henri Becquerel

• Discovered radioactivity by observing the spontaneous emission of radiation by uranium

  • Observed that a mineral containing uranium produces its image on a photographic plate in the absence of light

radioactivity

Gamma Rays

y symbol, high energy light

Beta Particles

β symbol, high speed electrons

alpha particles

α symbol, Particles with 2+ charge, Mass is 7300 times that of the electron

Gold Foil experiment

• done by Rutherford

• Carried out to test the accuracy of Thomson’s plum pudding model 

• Involved directing α particles at a thin sheet of metal foil 

• Expectation

  • α particles will pass through the foil with minor deflections in their paths

• Most α particles passed through the foil 

  • Atom is mostly open space 

• Many particles were deflected at large angles

  • Including those that had a close encounter with the massive positive center of the atom

• Some particles were reflected 

  • Including those that made a direct hit on the massive positive center

Nuclear atom

Has a dense center of positive charge called the nucleus with electrons moving around the nucleus at a distance that is large relative to the nuclear radius

Atomic structure

nucleus with protons and electrons, most of the mass, and electrons orbit nucleus

Protons

Have a positive charge that is equal in magnitude to the electron’s negative charge

Neutrons

 Have virtually the same mass as a proton but no charge

Electron mass and charge

Mass: 9.109 × 10^-31 kg

Charge: 1-

Proton Mass and Charge

Mass: 1.673 × 10^-27 kg

Charge: 1+

Neutron Mass and Charge

Mass: 1.675 × 10^-27 kg

Charge: 0

Isotopes

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

• Depict almost identical chemical properties

• In nature, most elements contain mixtures of isotopes

Atomic number

Number of protons, written as subscript

Mass Number

total number of protons and neutrons, written as superscript

Average diameter of atom

2 × 10^-10

Covalent bond

formed by sharing electrons

molecule

two or more atoms held together by chemical bond (can be the same element)

Chemical formula

CO₂ etc.

Structural Formula

• Depicts individual bonds in a molecule 

• May or may not indicate the actual shape of the molecule

Space-filling model

Illustrates the relative sizes of atoms and their relative orientation in the molecule

Ball and Stick Model

way of representing atoms, name is self explanatory

Ion

atom/group of atoms with a net charge

Cation

Positive ion

Anion

negative atom

Ionic Bonding

Force of attraction between oppositely charged ions

ionic solids

Solids containing oppositely charged ions

polyatomic Ions

have many atoms

Metals

• Efficient conductors of heat and electricity, malleable, and ductile

• Have a lustrous appearance 

• Tend to lose electrons to form positive ions

Nonmetals

• Lack the physical properties that characterize metals 

• Tend to gain electrons in reactions with metals to form negative ions 

• Often bond to each other by forming covalent bonds

Alkali Metals

• Members of Group 1A

• Very active elements that readily form ions with a 1+ charge when they react with nonmetals

Groups/Families

Elements in the vertical columns with similar chemical properties

Alkaline earth metals

• Members of Group 2A

• Form ions with a 2+ charge when they react with nonmetals

Halogens

Group 7a

• Form diatomic molecules

• React with metals to form salts containing ions with a 1– charge (exception - Astatine)

Noble Gases

• Members of Group 8A 

  • Exist under normal conditions as monatomic gases

  • Have little chemical reactivity

Periods

horizontal rows

Binary ionic compounds

Contain a cation, which is written first in the formula, and an anion (metal and non-metal)

Naming Binary Ionic compounds (type I)

• Cation is always named first and the anion second

• Monatomic cation takes its name from the name of the parent element

• Monatomic anion is named by taking the root of the element name and adding -ide

• No transition metals

Naming Binary Ionic Compounds (type II)

• Nomenclature for metals that form more than one type of cation 

  • Charge of the metal cation is indicated by a Roman numeral

• Alternative nomenclature 

  • Ion with the higher charge has a name ending in -ic, and the one with the lower charge has a name ending in -ous

Oxyanions

• type of poly atomic ions

• Anions that contain an atom of a given element and different numbers of O atoms

  • When there are two members in the series:

    • Name of the member with the smaller number of O atoms ends with -ite

    • Name of the member with the larger number of O atoms ends with -ate

• When more than two oxyanions make up a series:

  • Use the prefix hypo- (less than) to name members of the series with the fewest O atoms 

  • Use the prefix per- (more than) to name members of the series with the most O atoms

Binary Covalent compounds

two non-metals

Naming Binary Covalent Compounds

• First element in the formula is named first, using the full element name

• Second element is named as if it were an anion

• Prefixes are used to denote the numbers of atoms present

• Prefix mono- is never used for naming the first element

1 (prefix)

mono-

2 (prefix)

di-

3 (prefix)

tri-

4 (prefix)

tetra-

5 (prefix)

penta-

6 (prefix)

hexa-

7 (prefix)

hepta-

8 (prefix)

octa