chem test unit 4 - periodic table

Döbereiner

• Döbereiner triads

• Middle of the element is the average of the two outside masses

Mendeleev

• First periodic table organized by mass

• Patterns emerged but there was some problems like Te+I

Mosely

• Discovered the atomic #

• Took the periodic table and reorganized it by atomic number

• Invented the Periodic Law

Periodic Law

The properties of elements are in periodic function of the atomic #

metals

• good conductor of heat + electricity

• malleable

• luster (shiny)

• ductile

non-metal

• bad conductor

• exists in all states of matter

• brittle

• dull

noble gas

• group 18

• gases

• non-reactive “inert”

• “stable octet” (8 valence electrons)

metalloid

• mix of metallic and non-metallic

• properties depend on condition

• semi-conductor

metalloids found on the period table are

B, Si, Te, Sb, Ge, As, At

when metals bond with non-metals they

lose valence electrons + become positive

the easier you lose an electron

the more metallic you are

most metallic is

francium

most non-metallic is

flourine

what elements are gases at STP

O2 N2 Cl2 F2 H2 (ONCL Frank Has noble gas)

what elements are liquids at STP

Br2 Hg

what elements are solids at STP

every other element

monatomic gases

noble gases

diatomic

H2 N2 O2 Cl2 F2

allotrophs

are the same element in the same phase but have different molecular structures so the properties are different

ex. diamonds + graphite and oxygen + ozone

group 1

alkali metals

group 2

alkaline earth metals

group 3-12

transition metals

group 17

halogens

group 18

noble gas

atomic radius

½ distance between the nucleus of an atom and the outermost shell

ionic radius

the distance between the nucleus of an ion and the outermost shell

electronegativity

the ability of an atom to attract electrons

ionization energy

the amount of energy needed to remove electrons

effective nuclear charge

the net positive charge experience by valence electrons

what determines the properties of an element

• elements in a group have the same properties

• they have the same # of valence electrons

2 things that affect radius

• effective nuclear charge

• occupied principal energy levels

trends down a group - atomic radius

radius increases because the # of OPELS increases while the ENC stays the same

trends down a group - ionization energy

ionization energy decreases because as the radius increases the valence electrons are further away so it takes less energy to remove and electron

trends down a group - electronegativity

attraction for electrons decreases because as the radius increases the electron is further from the nucleus so it’s less attracted

trends down a group - metallic properties

increases because as the radius increases the electrons are further from the nucleus so they can be lost more easily

what affects effected nuclear charge

• atomic #

• shield electrons

trends across a period - atomci radius

decreases because the effective nuclear charge increases but it has the same OPELS

trends across a period - ionization energy

increases because as the radius decreases electrons are closer to the nucleus so it takes more energy to remove a valence electron

trends across a period - electronegativity

attraction for electrons increases because as the radius decreases the electron is closer to the nucleus making it more attracted

trends across a period - metallic properties

decreases because the electrons is more attracted to the nucleus so it’d be harder to remove

metal ionic radius

• lose electrons

• become positive

• smaller radius

non-metal ionic radius

• gain electron

• become negative

• radius increases

metal ionic radius reasoning

• the ion could lose and OPEL

• same # of protons but less electrons so it takes up less space

non-metal ionic radius reasoning

• adding an electron to an already occupied energy level increases repulsion so it will expand

transition metals always

lose higher energy level not sublevel

transition metal chemical properties

• less reactive than groups 1 + 2

• they have multiple oxidation levels

transition metals physical properties

form colored ions (1+2+12 form white)

why do they form colored ions?

split d-level

• all orbitals in the 3d don’t have the same energy

• so when the electrons move from a lower d-orbital to a higher one they absorb a photon of light which makes them appear colored

anomaly in ionization energy scenario 1

the valence electron is at a higher energy level because it’s the first in the highest sublevel so it takes less energy to remove (ex. Be and B)

anomaly in ionization energy scenario 2

the atom has a paired electron which experiences repulsion so it takes less energy to remove (N and O)