HONORS CHEMISTRY CHAPTER 3

Atoms: The Building Blocks of Matter

ICONS OFEARLY ATOMIC

THEORY

PART 1

ICONS IN EARLY ATOMIC THEORY Democritus [400 B.C]

Greek philosopherHypothesized: Nature

has a basic indivisible particle of which everything is made of Called this particle an

atomGreek “atomos” =

indivisible

1790S – DISCOVERY OF BASIC LAWS Law of Conservation of Mass

Mass is neither created nor destroyed during ordinary chemical reactions or physical changes

Law of Definite ProportionsA chemical compound contains the same

elements in exactly the same proportions by mass regardless of size of sample or source of compound i.e. Every sample of table salt is made of

39.34% Na and 60.66% Cl i.e. H2O always has 2 atoms of H and 1 atom of

O

BASIC LAWS CONTINUED Law of Multiple Proportions

If two or more different compounds are composed of the same two elements then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers i.e. CO and CO2

CO = 1.00g of C and 1.33 g of OCO2 = 1.00 g of C and 2.66 g of OThe ratio of the second element is 2.66 to 1.33 or 2

to 1

ICONS OFEARLY ATOMIC THEORY CONTINUED

John Dalton [1808]English schoolteacher – liked nature and

weatherDeveloped: Dalton’s Atomic Theory

DALTON’S ATOMIC THEORY1. All matter is composed of extremely

small particles called atoms2. Atoms of a given element are identical

in size, mass and other properties and are different from atoms of other elements

3. Atoms cannot be subdivided, created, or destroyed

4. Atoms of different elements combine in simple whole number ratios to form chemical compounds

5. In chemical reactions, atoms are combined, separated or rearranged

ISSUES WITHDALTON’S ATOMIC THEORY

Atoms can be split into even smaller particles (nuclear chemistry) and aren’t indivisible i.e. nucleus, protons, electrons

A given element can have different masses i.e. isotopes

STRUCTURE OF THE ATOM Today’s definition of the atom

Atom = Smallest particle of an element that retains the chemical properties of that element Two regions

Nucleus Very dense, small center of the

atoms Protons and neutrons

Electron Cloud Region occupied by electrons

Subatomic particlesProtons, neutrons, electrons

ICONS OF EARLY ATOMIC THEORY CONTINUED

J.J. Thomson [1897]Discovered: The 1st

subatomic particle: the negatively charged electron

Used a Cathode Ray Experiment Cathode Ray Tube –

Electric current passed through a metal disk to another metal disk in a gas at low pressure (vacuum sealed tube)

i. e. neon signs and ‘old-fashioned’ television sets

CATHODE RAY EXPERIMENT When a current passed

through the cathode ray tube, the surface of the tube opposite the cathode glowed Glow was hypothesized to be

stream of particles called a cathode ray

Ray affected by magnetic fields Attracted to positive

charge Deflected from negative

charge http://www.youtube.com/watch?

v=7YHwMWcxeX8&NR=1

DISCOVERY OF THE1ST SUBATOMIC PARTICLE Thomson measured the ratio of the

charge of the particles to their massSame ratio no matter what metal or gas

was usedNamed this particle an electron

http://www.youtube.com/watch?v=IdTxGJjA4Jw&feature=related

THOMSON’SPLUM PUDDING MODEL Atoms are electrically neutral

Must have positive charges to balance the negatively charged electrons

Electrons have a lot less mass than atomsOther particles must account for their mass

Plum Pudding Modelpositively charged sphere with electrons

dispersed through it

ICONS OFEARLY ATOMIC THEORY CONTINUED

Robert Millikan [1909]Discovered: The

measurement of an electron charge

Oil Drop Experiment Measured the

difference in velocity of oil dropletsCharged droplets

(ionizing radiation) vs. uncharged

http://www.youtube.com/watch?v=XMfYHag7Liw&feature=related

ICONS OFEARLY ATOMIC THEORY CONTINUED Ernest Rutherford (with Hans Geiger and

Ernest Marsden) [1911]Discovered: A new atomic modelGold Foil Experiment

Bombarded thin piece of gold foil with alpha particles Expected alpha particles to pass through with minimal

deflection Surprised when 1 in 8000 deflected back to source

It was “as if you had fired a 15 inch artillery shell at a piece of tissue paper and it came back and hit you”

http://www.youtube.com/watch?v=wzALbzTdnc8&feature=related

RUTHERFORD’SNEW MODEL OF THE ATOM Discovered the nucleus is a small

densely packed volume of positive chargeSize comparison

Nucleus = marble Whole Atom = football field

At this point in history, we were not sure where the electrons were – stay tuned for more in Chapter 4

INSIDE THE ATOM

PART 2

INSIDE THE NUCLEUS 2 types of particles

Protons positively charged = +1 made up of quarks

Neutrons neutral = 0 charge Made up of quarks

Mass in the nucleus Protons = 1.673 x 10-27

Neutrons = 1.675 x 10-27

To simplify, both have mass of 1 amu (atomic mass unit)

HOW DOES THE NUCLEUS STAY TOGETHER?

Strong Nuclear Forces Two protons extremely

close = strong attraction Two neutrons extremely close = strong attraction

Neutrons and Protons extremely close = strong attraction

Strong nuclear forces overcome the repulsion of like positive charges to keep the nucleus together!!!

WHERE ARE THE ELECTRONS? In the Electron Cloud

A cloud of negative charge outside of the nucleus

More on this later........ Electrons = Negatively charged particles

with almost no mass (9.109 x 10-31)

CHARACTERISTICS OF ATOMS Atomic Number

Equal to the number of protons and specific to each type of element

Identifies the element # of protons is what give that element its

characteristic properties Elements with different protons are NOT THE SAME

ELEMENT!!!

NEUTRAL ATOMS Neutral atoms

total positive charge equals the total negative charge # protons (+1 each) = # electrons (-1

each)

ISOTOPES Atoms of the same element (i.e.

same # of protons) that have differing number of neutrons

Isotopes of the same elementhave different massesdo not differ significantly in chemical

behavior

MASS NUMBER Mass number = #protons + #

neutrons

Average Atomic Mass Every element has isotopes The periodic table takes into account

all naturally occurring isotopes of an element and averages them

Element

Atomic Number

# of Protons

# of Neutrons

Mass Number

Carbon 6 6

8 16

Nitrogen

15

PROPERTIES OF SUBATOMIC PARTICLESParticle Symbol Charge Mass

Number

Electron e-, 0e -1 0

Proton p+, 1H +1 1

Neutron n◦, 1n 0 1

IONS Atoms with a charge

Negative – more electrons than protonsPositive – more protons than electrons

Charge = #protons - # electrons Magnesium atom with 12 protons and

10 electrons has a charge of +2

AVERAGE ATOMIC MASS Average Atomic Mass listed on the

periodic table UNIT is amu = atomic mass unit

1 amu is a standard Equal to 1/12 the mass of a C-12 atom

Takes into account all an elements isotopes and the frequency of each isotopes occurrence in natureHow to Calculate Average Atomic Mass

Mass of isotope #1 x abundance

in nature (decimal)

+ Mass of isotope #2 x abundance

in nature (decimal)

+ … =

Average Atomic Mass

EXAMPLE OF CALCULATING THE AVERAGE ATOMIC MASS – HYDROGEN

There are two naturally occurring isotopes of hydrogenHydrogen with 1 proton and zero neutronsHydrogen with 1 proton and one neutron

Differentiating between the two isotopes (symbol – mass number)

Calculation:

Hydrogen Isotopes

Element – mass #

Atomic Mass

Naturally occurring

abundance %

1 proton+ 0 neutrons

H-1 1.007825 amu

99.9885

1 proton+ 1 neutron

H-2 2.014102 amu

0.0115

PART 3: THE MOLE

AMADEO AVOGADRO Amadeo Avogadro [1776]

Lawyer turned professor of mathematical physics

Theorized: equal volumes of all gases at the same temperature and pressure contain the same number of particles.

After Avogadro’s death Avogadro’s number was determined

Avogadro’s number is simply a unit of measure 1 mole = 6.023 x 1023 of any substance

Typically used to talk about particles (atoms, compounds, etc.)

THE MOLE UNITPUT INTO PERSPECTIVE!!! One mole of rice grains is more grains

than the total number of grains grown since the beginning of time.

A mole of rice would occupy a cube about 120 miles on each edge.

A mole of marshmallows would cover the US to a depth of 600 miles

A mole of hockey pucks would be equal in volume to the moon

A mole of basketballs would just about fit perfectly into a ball bag the size of the earth.

MOLE VIDEO http://www.youtube.com/watch?v=Hj83

oRHdezc

MOLE CALCULATIONS 1 mole = 6.02 x 1023 of anything (atoms,

molecules, formula units, particles, etc.) Use dimensional analysis when solving:

Conversion factor:1 mole = 6.02 x 1023 atoms, particles, formula units, etc.

Practice:A. If I have 3.5 moles of carbon atoms, how

many molecules do I have?

B. If I have 5.43 x 1031 molecules of carbon dioxide, how many moles do I have?

MOLAR MASS Molar Mass

The mass of one mole of a substance The molar mass of an element canbe found on the periodic table

Same as the average atomic mass1 amu = 1 gram/mole

E.g. Average atomic mass of C = 12.011 amus Molar mass of C = 12.011 grams/1 mole

CALCULATIONS USING THE MOLAR MASS – USE DIMENSIONAL ANALYSIS

Calculate the number of grams of carbon in 3.25 moles of carbon.

Calculate the number of moles of hydrogen in 6.05 grams of hydrogen.

Calculate the number of atoms of carbon in 15.00 grams of carbon.

MOLE SING A LONG

October 24, 2005

IT’S A UNIT AFTER ALL”HTTP://WWW.YOUTUBE.COM/WATCH?V=ZNNBZGNSOHK

A mole of laughter, a mole of tears A mole of atoms, a mole of cheer The name of that measure Is a real chemist’s treasure It’s a unit after all Chorus

It’s a unit, after all It’s a unit, after all It’s a unit after all It’s a unit after all

IT’S A UNIT AFTER ALL” http://www.youtube.com/watch?

v=at_9A_gfln0 A chemist’s friend, tried and ture, An Avogadro would stand by you. And any chemist anywhere, Would stand up and swear, It’s a unit after all Chorus

It’s a unit after all, etc.

“IT’S A UNIT AFTER ALL” Six point oh two times ten to twenty-

three A number to live by in chemistry So this is October 24th Don’t be absurd, for It’s a unit after all Chorus

It’s a unit after all It’s a unit after all It’s a unit after all I’ts a unit after all