atomic theory

what did ancient greeks think about elements

4 elements - water, air, earth, fire

Democritus

thought substances were made of small particles, “atoms”

Aristotle

further developed Democritus’s idea

John dalton

• experimented with oxygen and other stuff

• tried to explain various compounds

dalton’s observations

• all matter made of atoms, small spheres

• matter cannot be created/destroyed/divided

• all atoms of element are identical

• different atoms have different properties

• each compound is unique, particular combination of atoms

• chemical reactions involve rearranging atoms

dalton’s laws

• law of definite proportions

  • compound contains fixed ratio of atoms

• law of multiple proportions

• law of conservation of mass

J.J. thomson

• discovered electron, plum pudding model

• cathode ray tubes, streams of negative particles

• theorized particles embedded in positive medium like pudding

Ernest Rutherford

• gold foil experiment, discovered nucleus

• fired alpha particles at gold foil, showed small concentrated positively charged region

• theorized positive nucleus surrounded by electrons buzzing like bees

Niels bohr

• electrons orbit nucleus in specific energy levels

• tried to explain why different elements have unique line spectra

• theorized electrons in atom were in specific energy levels

isotope

element with same atomic number but different masses

atomic number

protons in nucleus

atomic mass

protons + nucleus

neutrons

mass - protons

periodic table arranged according to

atomic number

group 1

• alkali

• soft, low density, highly reactive

group 2

• alkaline

• soft, low density, reactive

group 17

• halogens

• highly reactive non metals

group 18

• noble gases

• unreactive, used for lighting

where is atomic radius biggest

bottom left corner

where is atomic radius smallest

top right corner

where is first ionization energy greatest

top right corner

where is first ionization energy smallest

bottom left corner

electron configuration negative ions

add electrons to last unfilled orbital

electron configuration positive ions

lost electrons from orbital with highest coefficient, s before p before d

valence electrons

electrons involved in bonding

intramolecular forces

within a molecule, strong, ionic/covalent/metallic

intermolecular forces

between molecules or atoms, weak, hydrogen/dipole dipole/London, affect state/melting point

ionic bonding

metal/non metal, gives electrons

covalent bonding

• nonmetal bonding

• shares electrons

• have lone pairs

sigma bonds

• p orbitals overlap end to end

• electrons distributed along axis

• very strong

• occurs before pi bond

pi bonds

• p orbitals overlap side by side

• electrons distributed above/below bond axis

• weaker, occur after sigma bonds

• form when atoms make double/triple bonds

bond length

• distance between 2 nuclei of bonding atoms

• position themselves where lowest possible potential energy

• determined using X-ray crystallography

bond strength

• amount of energy required to break bond

• higher bond energy = stronger bond

• multiple bonds = stronger than single

electronegativity + bond polarity

• atom with greater electronegativity pulls electrons, becomes slightly negative

• atom with lower electronegativity becomes slightly positive

molecular polarity

• polar if have polar bonds and molecule not symmetrical

metallic bonding

• between metal atoms

• valence electrons loosely held, move freely around nuclei

• gives metal its properties

dipole

separation of charge, makes one end of particle slightly positive and other end slightly negative

electronegativity

tendency of atom to draw electrons towards itself

where are the most electronegative elements

top right corner of table

where are least electronegative elements

bottom left of table

London forces

• all molecules have

• caused by temporary dipole in non polar molecule or dipoles in polar

• affect clumping, so freezing

dipole dipole

• two elements with different electronegativities bond

• polar molecules only

• medium strength

hydrogen bonding

• strong

• only when H bonded to N, O, F

• other atom pulls electrons, H becomes positive

• explains why water is liquid, causes surfaces tension

valence electron shell pair repulsion

• shape of molecule affected

• valence electrons try to arrange far away from others

elements that form most bonds

C, Si, S, O, N

elements that bond only once

H, F, Cl, Br, I

where is most metallic element

bottom left corner

2 elements, no lone pairs

180, linear

3 elements, no lone pairs

120, trigonal planar

4 elements, no lone pairs

109.5, tetrahedral

5 elements, no lone pairs

90/120, trigonal bipyramidal

6 elements, no lone pairs

90, octahedral

1 element, 1 lone pair

n/a, linear

2 elements, 1 lone pair

120, bent

3 elements, 1 lone pair

109.5, trigonal pyramidal

4 elements, 1 lone pair

90/120, seesaw

5 elements, 1 lone pair

90, square pyramidal

1 element, 2 lone pairs

n/a, linear

2 elements, 2 lone pairs

109.5, bent

3 elements, 2 lone pairs

90, t-shaped

4 elements, 2 lone pairs

90, square planar

1 element, 3 lone pairs

n/a, linear

2 elements, 3 lone pairs

180, linear

3 elements, 2 lone pairs

90, t-shaped