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Chapter 8 - Atoms and Periodic Properties. Will turn to a study of the properties of matter why materials have certain properties chemistry - composition, structure and properties of substances and the transformations they undergo .

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slide1

Chapter 8 - Atoms and

Periodic Properties

Will turn to a study of the properties of matter

why materials have certain properties

chemistry - composition, structure and

properties of substances and the

transformations they undergo

consider world - many objects with many properties

trees bark, leaves, wood

car wheels, dash, hood

all substances made of combinations only

about 113 known elements

element - pure substance that cannot

be decomposed into simpler substances

by a chemical or physical process

well-defined properties

Water - H2O

Salt - NaCl

slide2

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Periodic Table - lists all known elements

currently about 108 known elements

88 naturally occurring

others made in lab

Some substances known for long time

how to describe?

alchemists - “lead into gold”

antimony Sb

confusing

STANDARDIZATION - modern symbols

used world-wide now

how did we get these symbols?

H He

C Cl

B Be

O Os

P Pt

S Se

N Ni

Hg

Ag

Au

Na

Fe

Sn

Pb

slide3

Where do the names come from?

Pu Es Am

U Fm Fr

Hg Md Eu

Np Bh Cf

Cl - light green

Tc - artificial

Ne - new

He - sun

Te - Earth

Elements named after planets, people,

places and descriptions!

Names passed by international council

“Commission on the Nomenclature

of Inorganic Chemistry”

Names agreed upon worldwide

standardized

slide4

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Elements made up of large collection of atoms

Atom - smallest unit with the same chemical

identity as element (10-10 m, 10-24 g)

chemical identity - physical and chemical

properties of a pure substance

structure of the atom:

protons, neutrons, and electrons

NUCLEUS - fixed central

part of atom

contains:

proton

positive charge

strong nuclear force

neutron

no charge

same mass as proton

does not affect

chemical identity

these are also called

nucleons-reside in nucleus

electron - negative charge (same as proton)

-swarm around nucleus (electron cloud)

-can be attracted away or added w/o chemical change

-very light 1/1837 mass of proton (negligible)

neutral atom-same number of p+ as e- (zero net charge)

remember ion: atom with net charge

slide5

How to determine atomic structure

History: Ancient Greeks

Democritus-matter is discontinuous

cannot divide indefinitely

“atom” - Greek for uncuttable

Aristotle & Plato disagreed with this view

(wrongly) thought matter was continuous

John Dalton (1800’s) revisited idea of Atoms

Dalton’s Atomic Theory

  • All matter = indivisible atoms
  • An element is made up of identical atoms
  • Different elements have atoms with different masses
  • Chemical compounds are made of atoms in specific integer ratios
  • Atoms are neither created nor destroyed in chemical reactions
slide6

MODERN IDEAS – discoveries leading to atomic structure

– indirect observations

J.J Thompson (late 1800’s) – discovery of electrons

cathode ray tubes – eject particles from plates

-cathode rays found to be negative

(opposites attract-not light)

-deflect in magnetic field (current-moving charge)

- measured charge-to-mass ratio

(crossed electric&magnetic fields)

Robert Millikan (1906)

Oil drop experiment

-charged oil drops in

electric field

-electric force opposed

gravity – drop floats

-droplet charge in multiples

of electron chargeqe=1.6x10-19 C

-found electron mass by using q/m from Millikan

me=9.11x10-31 kg very very small

slide7

Early model of the atom

Plum pudding model

Electrons embedded in blob of positively charged matter

like “raisins in plum pudding”

But what is the positive charge that cancels tiny electrons?

Rutherford –

alpha particle

positive helium nucleus

scattering- shoot alpha

particle at gold sheet

Result :

-most of the alpha

particles passed through sheet

-some alpha particles back-scattered

Conclusion:

-atom contains small central part

most of mass nucleus

-electrons orbit at distance 100,000

times the size of the nucleus

the atom is mostly made

up of EMPTY SPACE

Nucleus later found to be made of

protons (Rutherford split nucleus)

and neutrons (Chadwick-1932)

slide8

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Describing the Modern Atom

atomic number – number of protons in nucleus

-describes identity of element

-neutral atom number of e- = number of p+

mass number – number of protons and neutrons in nucleus

indicates mass since the electrons are negligible

new mass scale – STANDARDIZE

atomic mass units (amu, dalton)

1 amu is about mass of a proton

amu defined by mass of carbon-12:

carbon 12: 6 protons and 6 neutrons

define to have mass of exactly 12 amu

ATOMIC MASS STANDARD

But mass number does not define element

can have different numbers of neutrons

For example: Lithium

Li

ATOMIC NUMBER: 3 3 3

MASS NUMBER: 3 amu 4 amu 5 amu

Isotope: elements with the same number of protons,

but different numbers of neutron – different mass numbers

slide9

How do we study isotopes?

Mass Spectrometer

Curve of ions depends on the

charge-to-mass ratio

-isotopes have different masses

N

S

Each isotope will form a spot

At different places on the screen

Ions accelerated

in electric field

oven

Natural Abundance-what percentage of each isotope

exists for each element

Mass number –refers a particular isotope - specific atoms

Atomic Weight (Mass) – weighted average of the masses for

different isotopes in a sample of an element

for the element in general (all isotopes)

Notation: describes atomic structure: for an isotope

number of protons, neutrons and electrons

mass number

Means atomic number=78

and mass number =112 amu

11278Pt

Atomic structure-

protons: 78

electrons: 78 neutral

neutrons: 112-78=34

atomic number

Example:

3517Cl

slide10

An important isotope : H lightest element

H D T

hydrogen deuterium tritium

11H 21H 31H

normal heavy radioactive

hydrogen hydrogen

hydrogen-1 hydrogen-2 hydrogen-3

Natural 99.98% 0.015% < 0.005%

Abundance

atomic weight: 1.008 amu

Remember atomic model: SOLAR SYSTEM MODEL

massive nucleus surrounded by electrons

problem: electron circles atom -

centripetal acceleration

classical charge radiates if accelerated

loses energye-falls into nucleus

New Theory needed -- F=ma didn’t work

Planck & Einstein :

matter absorbs discrete amounts of energy

QUANTA

slide11

BOHR MODEL: tried to match experiments involving

absorption and emmision of light from hot solids

and gases - line spectra

not derived but phenomenological

Bohr’s Theory:

1. Electrons orbit the nucleus at

specific distances from the nucleus

-allowed orbits

2. Electrons in allowed orbits

do not radiate energy

-contrary to classical theory

3. Electrons gain energy by “jumping” to

a higher energy (further) orbit

-lose energy by falling to a lower energy

-energy loss or gain in

the form of a photon- particle of light

“Qnantum Leap”

n=1

n=2

n=3

Explained line spectra - electrons in matter

gain (absorb) or lose (emit) photons to make

quantum leaps

slide12

Wave-Particle duality : light travels like

particles and waves

de Broglie : matter also travels like waves

electrons travel like waves

-normal objects have very small wavelength

-electron motion governed by wave properties

STANDING WAVE

SOLUTION

Only certain

wavelengths (energies)

will fit correctly around nucleus

ALLOWED ORBITS

Led to the development of

QUANTUM MECHANIC THEORY

Schrodinger Equation-solve with linear algebra

and differential equations

Solution: electron orbital - 3D region surrounding

nucleus where there is the greatest probability of

finding an electron

slide13

Consequences of quantum mechanics

Solution gives energy levels of electrons

surrounding nucleus

-gives electron configuration

-the arrangement of electrons in

orbitals and suborbitals about the

nucleus of an atom

-describes properties of atom

“fingers of the atom”

interact through electrons

PROBABILITY DENSITY - probality at a

particular position

Cannot isolate position of an electron

HEISENBERG UNCERTAINTY PRINCIPLE

cannot measure momentum (motion)

and position of electron exactly

SOLUTION TO WAVE EQUATION GIVES

QUANTUM NUMBERS

-describe energies of the electrons

-determine properties of electrons in atom

-gives framework to “build” atoms

-similar electron configuration gives similar properties

-restrictions on what numbers can be

slide14

n=1

n=2

n=3

QUANTUM NUMBERS - n,m,l,s

describe energy levels of electrons

Principal Quantum Number ( n )

main energy level of electron

-describes orbit

-how far electron is

from nucleus

-similar to allowed

Bohr orbit

-restriction: whole number

greater than 0

n=1,2,3,4,…

closest furthest

Angular Momentum Quantum Number ( l )

shape of electron orbit

-how spread out the orbital is

-restriction:l= 0 to (n-1)

l = 1

hourglass

l = 0

sphere

slide15

Magnetic Quantum Number ( m )

orientation of electron orbital

-the way the electrons are oriented about nucleus

-restriction:-l > m > +l

Example : n=2 electrons

l=0 electrons have possible m=0

only one way

to orient sphere

l=1 electrons have possible m= -1, 0, +1

oriented in y-dir

m=-1

oriented in x-dir

m=+1

oriented in z-dir

m=0

n, l, m describe spatial properties of electron

how the electron cloud looks

each quantum number describes electron

with specific energy in nucleus!

slide16

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Spin Quantun Number (s)

magnetic properties of the electron

electron - electric charge

spin clockwise spin counterclockwise

spin up spin down

electron magnets interact with magnetic field

split into two beams in magnet

- restriction s = +1/.2, -1/2

ELECTRON CONFIGURATION determined by the

values of quantum numbers n, l, m, s

“fingers of the atom”

how the electrons interact with their environment

PHYSICAL AND CHEMICAL PROPERTIES

(CHEMICAL IDENTITY)

S

N

N

S

slide17

CAN NOW BUILD ATOMS

Need some rules before building electron configurations

Electrically neutral atoms - same number of e- as p+

have no net charge

Ground state atom - electrons occupy only the lowest

energy levels in atom

as opposed to excited state - electrons occupy higher

energy states

lower energies unoccupied

Will add electrons up to atomic number

-but how do we add electrons?

TWO WAYS TO ADD PARTICLES:

1. Put all electrons in lowest energy level

{n=1, l=0, m=0, s=1/2}

cannot do for e- only for BOSONS (photons)

2. Pauli Exclusion Principle: no two electrons

in the same atom can have the same set of

quantum numbers n, l, m, s

Law of nature for fermions (spin=1/2)

slide18

BUILD TABLE OF ELEMENTS

First note: for the spatial orbital

n, l, m describe position in space

two values of s for each

nlm combination

Start building - add electrons successively to

each lowest energy orbital

H : atomic number = 1 one electron

put electron in lowest energy

Electron configuration:

n=1, l=0, m=0, s=1/2

He : atomic number = 2 two electrons

Electron configuration:

n=1, l=0, m=0, s=1/2 still lowest

n=1, l=0, m=0, s=-1/2 -different atom

just change to s=-1/2 (next energy)

LI : atomic number = 3 three electrons

Electron configuration:

n=1, l=0, m=0, s=1/2

n=1, l=0, m=0, s=-1/2 n=1 full, next n=2

n=2, l=0, m=0, s=1/2 electron capacity

Be to Ne are filled by adding two e- to each n, l, m

Electron capacity - maximum number of electrons

that can be added to each orbital

slide19

Connections to Periodic Table

Note: row (period) determined by highest

principal quantum number

electron capacity met at end of row:

NOBLE GASES

outer shell not full - chemically reactive

outer shell full - no electrons to interact

with other elements

chemically inert NOBLE GAS

INERT GAS

get to next element by adding electron

to next level in orbital up to electron capacity

Electron properties determined by principle (n) and

angular momentum (l) Q.N

Electron orbital notation:

specify n, l, and number of electrons

in the orbital (superscript)

electron capacity

s orbital l =0 2

p orbital l =1 6 electrons

d orbital l =2 10 allowed in the

f orbital l =3 14 orbital

Example : 3d2 n=3, l =2, 2 e- in orbital

slide20

Rewrite electron configuration in new notation

H Li Na K

Be Mg Ca

B Al

C Si

N P

O S

F Cl

He Ne Ar

Sc next: things are different, but first

Periodic (Moseley’s) Law - electron configurations repeat

properties of elements repeat when

ordered by increasing atomic number

periodic function of atomic number

similar outer shell, similar properties

First column : Alkali metals s1 orbital

very reactive - single electron

Second column: Alkaline Earth Metals s2 orbital

Last column : (Inert) Noble Gases s2p6 (s2 for He)

very stable - octet (eight outer e- except He)

Electron configuration repeats, chemical properties repeat

slide21

Led to Periodic Table

Mendelev thought to be the father of the Periodic Table

Periodic Table – table of all know elements

listed in order of atomic number

-periodicity in properties along rows

(density, melting/boiling, hardness, etc)

DIVIDED INTO:

Families (Groups) – vertical column of elements

these elements exhibit similar properties

have same outer electron configuration

Eight main groups : Group IA to Group VIIIA

MAIN GROUP or REPRESENTATIVE GROUPS

show similarities in outer e- shell ( want octet )

Group IA – Alkali Metals (react violently w/H2O) s1

never uncombined in nature

Group IIA – Alkaline Earth Metals (also reactive) s2

Group IIIA s2p1

Group IVA s2p2

Group VA s2p3

Group VIA s2p4

Group VIIA –Halogens (salt former w/metal) s2p5

Group VIIIA - Noble (Rare)Gasses s2p6

never bond with others

{

explains

Periodic Law

slide22

Transition Metals – B groups

Group IB to Group VIIB

fill inner electron orbitals

like Sc: 1s22s22p63s23p64s23d1 skips energy

change in orbital energies

higher orbitals have lower energy

i.e., 4s is lower than 3d

shows gap in periodic table

electron energy level order:

1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p6f7d

increasing energy

can write any electron configuration

Example 1. Write configuration for:

Zr

V

Example 2. Identify element electron configuration

1s22s22p63s23p64s23d7

1s22s22p63s23p4

Periods – row groupings of the Periodic Table

properties repeat as you go from one period to the next

-periods begin reactive (IA) and end stable (VIIIA)

Note: Period and Group of element identified properties

Historically: some elements undiscovered-

chemists knew properties before

it even existed

slide23

How to read Periodic Table

Name Group II (family): 2 e-

Atomic number group: # of outer e-

Symbol

Atomic weight Period 4

Magnesium

12

Mg

24.31

Electron dot notation

KERNEL – nucleus and inner electrons

-dots represent the outermost electrons

-shows what’s available for the other

atoms to interact with

C group IV 4 outer electrons

Metals non-metals and semiconductors (semimetals)

METALS- conducts heat and electricity

- metallic luster (shiny)

- maleable pond into sheets

- ductile  draw into wires (extrusion)

-form positive ions by losing electrons

Li 1s22s1 Li+ stable [He] config

Mg 1s22s22p63s2  Mg+2stable [Ne] octet

NONMETALS- insulators

- dull appearance

- brittle

- form negative ions to complete octet

Cl 1s22s22p63s23p5+1 Cl-stable [Ar] config

X

X

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