young people should not smoke. young people should not smoke. studies show that smoking at an early...

• Young people should not smoke.

• Studies show that smoking at an early age may make it more difficult to quit smoking later.

• Which of the above statements is an opinion and which is a theory?

• What is the difference?

Atomic TheoryWhat is it?Who figured it out?When did they do it?How did they do it?Why do we believe it?

What can you tell from this picture?

What is an ATOM?

An atom is the smallest particle of an element that retains the identity of the element.

ZZZ

ZZZ

ZZZ

ATOM

Atoms combine to form compounds

The Ancient Greeks

DemocritusLived 450 B.C.Proposed that all

the stuff in the world is “atomos”

Tiny, indivisible particles

AristotleQuestioned theory

of DemocritusRejected it for lack

of proof

Neither offered PROOF, but Aristotle was

respected as the smartest guy around…the

“Einstein” of his time! So Democritus’ idea

was ignored!

ALCHEMISTS

Tried to turn base metals into precious metals.

Developed knowledge and techniques…very smart and knew a lot!

Not true scientists Didn’t share!

Roger Bacon

Lived in 13th centuryBelieved that science should be based on

experimental evidence.

Antoine Lavoisier 1743-1794 French Father of Chemistry Sadly, beheaded in French Revolution

Lavoisier

LAW OF CONSERVATION OF MATTERMatter, like energy, is neither created nor

destroyed in a chemical reaction.

This concept established modern chemistry.

Law of Conservation of Mass

What was different?

Experimentation He used a balance to

study the role of oxygen in rusting and burning.

Other people did it, too!

Priestly did similar experiments.

He believed a false theory of Phlogiston.

Kept his mind closed to new idea.

PROUST

1799

Law of Constant CompositionA given compound

always contains the same elements in the same proportion by mass.

OXYGEN89%

HYDROGEN11%

Law of Multiple Proportions(stated by Dalton)

The ratios of the masses of elements

in compounds will be ratios of small

whole numbers…multiple proportions.

JOHN DALTON

LAVOISIER PROUST

DEMOCRITUSBACON

ATOMIC THEORY OF MATTER

DALTON’S ATOMIC THEORY OF MATTER

1. All matter is composed of submicroscopic (extremely small) indivisible particles called ATOMS.

2. All atoms of a given element are identical. The atoms of any one element are different from those of any other element.

DALTON’S ATOMIC THEORY OF MATTER

3. Atoms are neither created nor destroyed in a chemical reaction. Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another element as a result of a chemical reaction.

4. Atoms of different elements can mix physically or can combine chemically with one another in simple whole number ratios to form compounds.

Dalton’s Theory

Based on

CAREFUL EXPERIMENTATION

DALTON’S MODEL OF THE ATOM

1803

Modern Atomic Theory

Not all aspects of Dalton’s atomic theory have proven to be correct. We now know that:--Atoms are divisible into even smaller particles.

--A given element can have atoms with different masses.

Some important concepts remain unchanged.

--All matter is composed of atoms.

--Atoms of any one element differ in properties from atoms of another element.

What does this mean?

Atomic Theory of Matter

Not perfect...

but a workable theory to build on.

Scientists had been doing many experiments with electricity since Ben Franklin flew his kite.

Faraday suggested that electricity might explain the atom

English physicist, J.J. Thomson 1856-1940

Discovered electrons in 1897

Thomson experimented with a “cathode ray tube”

Thomson’s Cathode Ray Tube Experiment

Thomson used ELECTRICITY To probe the atom

Cathode Ray Tube

Flow of electric current through gases.

Sealed gas in glass tube with metal plates at the end.

Connected plates to high voltage source:

Cathode – and Anode +

CRT

+

This ray could be deflected toward a positive charge…

It has a negative charge.

This ray could move things…

It was made of particles of matter.

This ray acted the same no matter what materials were used…

It was not atoms…it must be part of all atoms

Cathode Ray Tube

Cathode ray is composed of very small negatively charged particles that are part of

atoms

ELECTRONS

Thomson’s Plum Pudding Model

1896 Do you like

Chocolate chip

cookies better?

Chocolate Chip

Cookie Dough

Millikan’s Oil Drop Experiment1909

Determined the

charge of the

electron which

along with

Thomson’s

experiment

determined the

mass of the

electron

Atomic Theory

Millikan devised experiments to determine the mass and charge of the electron.

Protons : discovered in 1886

Positively charged particles

Also discovered with cathode ray tube…these particles went the other direction.

The Nuclear AtomWhat’s missing?

The Nucleus!Rutherford did an experiment using

RADIOACTIVITY as a tool to probe the atom

Identified three types of radiationAlpha…positive particleBeta…negative particleGamma…high energy

Devised the GOLD FOIL EXPERIMENT

A. The Structure of the Atom

Rutherford’s Experiment

(a) The results that the metal foil

experiment would have yielded if the

plum pudding model had been correct

(b) Actual results

Unexpected Results of the Rutherford Experiment

Rutherford’s Nuclear Atom1911

Most of the mass of an atom is in the center…the nucleus…with electrons moving around it.

Nuclear atom

By 1932, neutron was discovered, too.The nucleus is the central core of the

atom, composed of protons and neutrons. Because, protons and neutrons have much greater mass than electrons, almost all of the mass of the atom is concentrated in a tiny nucleus…a dime in a football stadium!

Forces in the Nucleus

• When two protons are extremely close to each other, there is a strong attraction

between them.

• A similar attraction exists when neutrons are very close to each other or when

protons and neutrons are very close together.

•The short-range proton-neutron, proton-proton, and

neutron-neutron forces that hold the nuclear particles

together are referred to as nuclear forces.

What next?

We know the parts of the atom.We know about the nucleus.What about the electrons?

Bohr’s Model

LIGHT provided the next clues for the structure of the atom

Light has a Dual Nature

Like this picture…

Young lady or

Old lady?

Light is a wave

Electromagnetic radiation travels in waves.

As the wavelength increases, the frequency decreases.

The greater the frequency the greater the energy.

Light acts like Particles

Electromagnetic radiation also has the properties of particles.

Bohr suggested that energy is emitted and absorbed in discrete quantities called

Quanta or Quantumpackets or pieces of energy

“Jumps”

Dual Nature of Light is the next tool for understanding the atom

Energy is directly proportional to frequency…wave nature.

Example: Light diffuses through small slitsEinstein proposed that light consists of

quanta of energy that behave like particles of light…he called these photons.

Example: photoelectric effect…garage door openers, film

This is the DUAL NATURE OF LIGHT

Demonstrate Quantum

Continuous Spectrum vs. Line Spectrum

Emission Spectra

Line Spectra of elements show discrete or quantized energy levels in the atom.

These energy levels are different for each element.

The color or wavelength of light shows the energy level because energy can be calculated from frequency of light emitted.

Quantum Leaps!

Electrons exist at low energy…ground state

Add energy to go to higher energy…excited state• Electron drops back down to ground

state as soon as it can • It releases the exact energy it needed to

jump up as a photon of a frequency or color.

Bohr’s Model of the Atom1913

Bohr applied Quantum Theory to the structure of the hydrogen atom.

Quantum theory means that electrons jump from level to level.

Bohr pictured the electrons spinning around set orbits, like planets around the sun.

Planetary Model

Schroedinger’s Quantum Mechanical

(Today’s) ModelWe still have energy levels, like Bohr

MATH is used as the tool to probe the atom

Heisenberg Uncertainty PrincipleIt is impossible to know the velocity and

position of a particle at the same timeWe can not know the exact location of an

electron.Any effort to do so, will change the position.We can only figure the probability of finding it

in a certain region.

Orbitals

Area where electrons are most likely to be found.

Example: If your mom wants you to do some work, she can find you in your room most of the time, or the house, or the yard…but you could be at the mall.And sometimes her efforts to find you, make you move!

Today’s Atom1926

Dense, small,nucleus (with

protons and neutrons) is

surrounded by a fuzzy cloud

shapes where electrons (that act

more like waves) are most likely to

be found… orbitals.

Today’s Model of the Atom

aka

Schroedinger’s

or

Quantum Mechanical Model

The Crazy World of Quantum Mechanics“If you aren’t shocked by it, you don’t

understand it.” Bohr

An atom is the smallest particle of an element that retains

the chemical identity of the element.

Three Parts of Atom

Particle Charge Mass Atomic mass units

Proton +1 1.67x 10-24 g 1 amu

Neutron 0 1.67x 10-24 g 1 amu

Electron -1 9.11x 10-28 g 0 amu

Subatomic particles

B. Introduction to the Modern Concept of Atomic Structure

Comparing the Parts of an Atom

Where are they?

Nucleus is made up of protons and neutrons.

Electrons move around the nucleus.Mass is concentrated in the nucleus.99.97% of the mass is in the nucleus.

HOW MUCH FOR A ROOT BEER?

A NEUTRON WALKED INTO A BAR AND ASKED…

FOR YOU, BUDDY…

NO CHARGE!

THE BARTENDER SAID

How big are atoms?

If nucleus is size of a golf ball, the electrons are how far away?

1 mile! Lots of empty spaceHow big are they?6.02 x 1023 atoms in 12 grams of carbon.

How many protons?

The atomic number is the number of protons in the nucleus of the atom.

Look at the periodic table for this.

The number of protons identifies the element.

CARBON

Carbon has 6 protons 6 amu6 neutrons 6 amu6 electrons 0 amu

Total atomic mass = 12 amu

How do we write this? Carbon-12 C-12 12C 12 6 C0

The periodic table tells us that carbon has 6 protons.

The mass is 12 amu and equals protons plus neutrons, so it has 6 neutrons.

It is neutral (has no charge), so it has 6 electrons.

How many neutrons?

The number of neutrons may vary. An isotope is an atom that has the same number

of protons as other atoms of the same element, but different number of neutrons.

Same atomic number = same element Protons + Neutrons = Atomic Mass Isotopes have different atomic masses, because

they have different number of neutrons.

Isotopes Isotopes are atoms with the same number

of protons but different numbers of neutrons.

Isotope of carbon Carbon-14 C-14 14C 14 6 C0

The periodic table tells us that carbon has 6 protons.

The mass is 14 amu and equals protons plus neutrons, so it has 8 neutrons.

It is neutral (has no charge), so it has 6 electrons.

What’s the safe answer

here?

???

Carbon-14 isotope said to Carbon-12

isotope,

“Do these neutrons make my mass look

big?

How many electrons?

The number of electrons equals the number of protons in a neutral atom

An ion is a charged atom in which the number of electrons can increase or decrease.

More electrons than protons, charge is negative

Fewer electrons than protons, charge is positive

Ion of carbon Carbon-12 C-12 12C 12 6 C+4

The periodic table tells us that carbon has 6 protons.

The mass is 12 amu and equals protons plus neutrons, so it has 6 neutrons.

It is positive (it lost 4 electrons), so it has only 2 electrons.

Are you POSITIVE?

OH NO!!!

I think LOST an

ELECTRON!!

A n (- or +)

E

Z

E = element symbol

Z = atomic number = # protons (identifies element)

A = atomic mass = # protons + # neutrons

n = charge

(+ if fewer electrons than protons)

(- if more electrons than protons)

AVERAGE ATOMIC MASSis a

WEIGHTED AVERAGE• Imagine that your semester grade depends 60%

on exam scores and 40% on laboratory explorations.

• Your exam scores would count more heavily toward your final grade.

• In this section, you will learn that the atomic mass of an element is a weighted average of the masses of the naturally occurring isotopes of that element.

Relative Atomic Masses• The standard used by scientists to compare

units of atomic mass is the carbon-12 atom, which has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu.

• One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom.

• The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

2 examples

Test scores for a class

2 examples

Easy atoms

Average Atomic Masses of Elements

• Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. It accounts for the relative abundance of each isotope

Calculating Average Atomic Mass• The average atomic mass of an element

depends on both the mass and the relative abundance of each of the element’s isotopes.

Calculating Average Atomic Mass

• Copper consists of 69.09% copper-63, which has an atomic mass of

62.9298 amu, and 30.91% copper-65, which has an atomic mass of

64.9278 amu.

• The average atomic mass of copper can be calculated by multiplying the

atomic mass of each isotope by its relative abundance (expressed in decimal

form) and adding the results.

Calculating Average Atomic Mass

• Copper consists of 69.09% copper-63, which has an atomic mass of

62.9298 amu, and 30.91% copper-65, which has an atomic mass of

64.9278 amu.

• The average atomic mass of copper can be calculated by multiplying the

atomic mass of each isotope by its relative abundance (expressed in decimal

form) and adding the results.

• (0.6909× 62.9298 amu) + (0.3091 × 64.9278 amu) =

63.55 amu

• The calculated average atomic mass of naturally occurring copper is

63.55 amu.

Solve

Calculate the average atomic mass for element X, the element that is a “goofy prisoner!” Then identify the element.

X-28 27.977 amu 92.21%

X-29 28.976 amu 4.70%

X-30 29.974 amu 3.09%

Counting Atoms to Determine Molar Mass

We need to be able to count atoms to determine molar mass

FIRST: See handout to practice counting ALL atoms

Subscripts…# to right and below multiply what goes before it Parentheses…atoms within act as a group, subscript applies to

all of the elements in parentheses Coefficients…multiply everything

SECOND: IGNORING COEFFICIENTS, add masses of all atoms to

determine molar mass of each substance. Use masses to the hundredths place.

VOLUME OF A GAS (L)

(in Liters at STP)

MASS (g)

(mass in grams)MOLES (mol)

NUMBER OF PARTICLES (#)

(atoms, ions, molecules, formula units)

Molar mass from PT

grams / mole

Molar volume of a gas

22.4 L of a gas / mole of gas (at STP)

Avogadro’s #

6.022 x 1023 particles / mole

THE MOLE MAP

Everything goes through the

MOLE.

Use mole conversion factors and

factor-label to solve

Relating Mass and Volume to Numbers of Atoms

The mole (mol) is the SI unit for the AMOUNTof a substance. It relates mass or volume of a gas (things we can easily measure ) to the # of particles (things too small to measure)

• The number of PARTICLES of a substance. A mole of anything contains as

many particles as there are atoms in exactly 12 g of carbon-12.

• “particles” can be atoms, ions, molecules, or formula units

• Avogadro’s number—6.022 1415 × 1023—is the number of particles in

exactly one mole of a pure substance.

Conversions with Avogadro’s Number

• Avogadro’s number can be used to find the number of atoms of an element from the amount in moles or to find the amount of an element in moles from the number of atoms.

• In these calculations, Avogadro’s number is expressed in units of atoms per mole.

Sample Problem AHow many moles of silver, Ag, are in 3.01 1023 atoms of silver?

VOLUME OF A GAS (L)

(in Liters at STP)

MASS (g)

(mass in grams)MOLES (mol)

NUMBER OF PARTICLES (#)

(atoms, ions, molecules, formula units)

Molar mass from PT

g/mole

22.4 L of a gas/ mole of a gas at STP

6.022 x 1023 particles / mole

THE MOLE MAP

Everything goes through the

MOLE.

Use mole conversion factors and

factor-label to solve

Sample Problem A Solution

Given: 3.01 × 1023 atoms of Ag Unknown: amount of Ag in moles Solution:

moles Ag

Ag atoms = moles AgAvogadro's number of Ag atoms

2323

1 mol Ag3.01 10 Ag atoms

6.022 10 Ag at

0.500

=

m

oms

ol Ag

Gram/Mole Conversions

4.00 g He2.00 mol He = 8.00 g He

1 mol He

• Chemists use molar mass as a conversion factor in chemical calculations.

• For example, the molar mass of helium is 4.00 g He/mol He.

• To find how many grams of helium there are in two moles of helium, multiply by the molar

mass.

Molar Mass• The mass of one mole of a pure substance is called the molar mass

of that substance.

• Molar mass is usually written in units of g/mol.

• The molar mass of an element is numerically equal to the atomic

mass of the element in atomic mass units.

Sample Problem BWhat is the mass in grams of 3.50 mol of the element copper, Cu?

VOLUME OF A GAS (L)

(in Liters at STP)

MASS (g)

(mass in grams)MOLES (mol)

NUMBER OF PARTICLES (#)

(atoms, ions, molecules, formula units)

Molar mass from PT

g/mole

22.4 L of a gas/ mole of a gas at STP

6.022 x 1023 particles / mole

THE MOLE MAP

Everything goes through the

MOLE.

Use mole conversion factors and

factor-label to solve

Sample Problem B Solution

Given: 3.50 mol CuUnknown: mass of Cu in gramsSolution: the mass of an element in

grams can be calculated by multiplying the amount of the element in moles by the element’s molar mass.

grams Cumoles Cu × = grams Cu

moles Cu

Sample Problem B Solution, continued

The molar mass of copper from the periodic table is rounded to 63.55 g/mol.

63.55 g Cu3.50 mol Cu × =

1 222

mol Cu g Cu

Sample Problem CA chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced?

VOLUME OF A GAS (L)

(in Liters at STP)

MASS (g)

(mass in grams)MOLES (mol)

NUMBER OF PARTICLES (#)

(atoms, ions, molecules, formula units)

Molar mass from PT

g/mole

22.4 L of a gas/ mole of a gas at STP

6.022 x 1023 particles / mole

THE MOLE MAP

Everything goes through the

MOLE.

Use mole conversion factors and

factor-label to solve

Sample Problem C Solution Given: 11.9 g Al Unknown: amount of Al in moles

Solution:

moles Al

grams Al = moles Algrams Al

1 mol Al11.9 g Al =

26.0.441

98 g Al mol Al

The molar mass of aluminum from the periodic table is rounded to 26.98

g/mol.

Sample Problem DWhat is the mass in grams of

1.20 108 atoms of copper, Cu?

VOLUME OF A GAS (L)

(in Liters at STP)

MASS (g)

(mass in grams)MOLES (mol)

NUMBER OF PARTICLES (#)

(atoms, ions, molecules, formula units)

Molar mass from PT

g/mole

22.4 L of a gas/ mole of a gas at STP

6.022 x 1023 particles / mole

THE MOLE MAP

Everything goes through the

MOLE.

Use mole conversion factors and

factor-label to solve

Sample Problem D Solution Given: 1.20 × 108 atoms of Cu Unknown: mass of Cu in grams Solution:

moles Cu grams Cu

Cu atoms = grams CuAvogadro's number of Cu atoms moles Cu

14

823

1 mol Cu 63.55 g Cu1.20 10 Cu atoms =

6.022 10 Cu atoms 1 mol Cu

1. 27 10 Cu g

–

The molar mass of copper from the periodic table is rounded to 63.55 g/mol.

The Mole and Volume

The mole also relates the amount of atoms to the volume of a gas at STP, Standard Temperature and Pressure.

One mole of gas at STP is 22.4 Liters

VOLUME OF A GAS (L)

(in Liters at STP)

MASS (g)

(mass in grams)MOLES (mol)

NUMBER OF PARTICLES (#)

(atoms, ions, molecules, formula units)

Molar mass from PT

g/mole

22.4 L of a gas/ mole of a gas at STP

6.022 x 1023 particles / mole

THE MOLE MAP

Everything goes through the

MOLE.

Use mole conversion factors and

factor-label to solve

Sample Problem E

How many moles of Helium are present in 44.8 L at STP? How many grams?

44.8 L He | 1 mole He = 2 moles He

22.4 L He

44.8 L He | 1 mole | 4.00 g He = 8 g He

22.4 L He | 1 mole He