I. Law vs. Theory

1) Scientific law = a generalization of scientific observations that describes what happens (does not explain)

2) Theory (model) = a set of assumptions used to explain observations and predict new observations

THEORY:a) Can never be truly proven

100% correct.b) Inevitably change (and must

sometimes be abandoned) as more information becomes available.

c) Considered successful when they explain observations and, more importantly, predict new observations.

Atomic Structure1) EARLY ATOMIC THEORY:

a)400 B.C- Democritus (Greek philosopher)• Coined the term “atom” (meaning indivisible) to describe the smallest particles of matter

• Did not experiment

2) JOHN DALTON (EARLY 1800’S):a) “Father of the

Modern Atomic Theory”

b) Dalton’s Atomic Theory helped explain measurable observations and successfully predict new observations (while stimulating additional research from Dalton and other scientists).

John Dalton

c)Assumptions of Dalton’s Atomic Theory (proposed by Dalton in 1803):

I.All matter is composed of atoms

II.All atoms of the same element are identical in size, mass, other properties. Atoms of different elements differ in size, mass, and other properties.

John DaltonIII.Atoms cannot be subdivided,

created, or destroyedIV.Atoms of different elements

combine in simple whole-number ratios to form chemical compounds.

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

John Daltond) Dalton used his Atomic Theory to

help correctly predict new observations leading to the “Law of Multiple Proportions”.

John Daltone)Law of Multiple Proportions = The masses of

one element that combine with a constant mass of another element to form more than one compound are in the ratio of small whole numbers• Carbon monoxide and carbon dioxide

contain oxygen in a 1:2 ratio.• 12 g of carbon reacts with 16 g of oxygen

to form carbon monoxide, CO.• 12 g of carbon will react with 32 g of

oxygen to form carbon dioxide, CO2.

John Dalton

f) Dalton’s Atomic Theory explained many observations and correctly predicted many additional observations. It DID NOT correctly predict all new observations (no theory ever does).

g) Dalton’s Atomic Theory has been modifiedto explain the new observations:

• Atoms of one element can have differentmasses (isotopes).

• Atoms can be subdivided (but not in a chemical reaction or physical change)…Nuclear reactions

• An atom of one element can be changed into an atom of another element (but not in a chemical reaction or physical change)……Nuclear reactions

John Dalton:g) As with all theories,

Atomic Theory has been changed and expanded over time (and will continue to change and expand) to explain new observations.

3) CATHODE-RAY TUBE EXPERIMENTS (late 1800’s):• Cathode-ray tube = A glass tube containing a gas at a

very low pressure which contains a negative electrode (cathode) and a positive electrode (anode). When high voltage is applied, a “cathode ray” passes from the cathode to the anode causing the low pressure gas in the tube to glow (different gases glow different colors).

3) CATHODE-RAY TUBE EXPERIMENTS (late 1800’s):

a) Early cathode-ray tube experiments:• Proposed that a cathode ray consists

of tiny particles with mass and they are negatively charged.

3) CATHODE-RAY TUBE EXPERIMENTS (late 1800’s):

b) J. J. Thomson –discovered theelectron

• Used early cathode-ray tube experiments and some of his own findings to support the hypothesis that electrons are negatively charged particles

J.J. THOMSON:• Used a cathode ray tube- as

voltage across the tube was increased, a beam of light (cathode ray) became visible

• Cathode ray = beam of electrons seen b/c of excited gas

J.J. THOMSON:• Found that the

beam was deflected by both magnetic and electrical fields

• OBSERVATION:Noticed that the “cathode rays” were attracted to the positive electrode, called the anode.

What conclusions did Thompson have?• The cathode rays were made up a very

smallnegatively charged particle –ELECTRONS

• OBSERVATION: Measured the bending of the path of the cathode rays and was able to determining the ratio of an electron’s charge to its mass. Found that the ratio was always the same, regardless of the metal used

J.J. THOMSON:

What conclusions did Thompson have?

• Concluded that the negatively charged particles were much lighter than the lightest know atom (hydrogen), which meant that atoms had a structure!

c) Later Protons were also discovered using a modified cathode-ray tube

d) Thomson’s Model of the Atom (Plum Pudding Model)• Thomson postulated that an atom consisted of a

diffuse cloud of positive charge with the negative electrons embedded randomly in it.

• This represented a major change from Dalton’s model of atoms as indivisible.

• The existence of the electron raised new questions: if electrons are part of all matter and they possess a negative charge, how can all matter be neutral? Also, if the mass of an electron is so small, what accounts for the rest of the mass in a typical atom?• More experiments to come…

4) ERNEST RUTHERFORD & THE GOLD FOIL

EXPERIMENT

a)Rutherford discovered the NUCLEUS using alpha particles & gold foil• Alpha particle = a

relatively large positively charged particle

ERNEST RUTHERFORDb) Gold-Foil Experiment (around 1908-1909):

Thin gold foil was bombarded by alpha particles and the path of the alpha particles was charted after they passed through the gold foil.

ERNEST RUTHERFORDc)Expected results: The

massive alpha particles (positively charged) were expected to “crash” through the gold foil with little or no deflections; there was nothing in Thomson’s model of the atom to cause anything more than minor deflections of the alpha particles.

ERNEST RUTHERFORDd)OBSERVATION: Some alpha particles

were deflected at large angles, and some were redirected backward (surprising results!)

• What conclusion did Rutherford draw from this evidence?The POSITIVE particles of the atom

must NOT be spread out evenly, but instead must be concentrated at the center of the atom- the (nucleus).

ERNEST RUTHERFORDe)OBSERVATION: Most of the alpha

particles passed through the gold foil with few deflections.• What conclusion did Rutherford draw

from this evidence?Most of the alpha particles did not hit anything and passed straight through the gold atoms so therefore, most of the volume of an atom consists of empty space.

ERNEST RUTHERFORDf) Rutherford’s Nuclear Model

• Explained the neutral nature of matter: the positive charge of the nucleus balances the negative charge of the electrons.

ERNST RUTHERFORD• Suggested that the electrons travel around

the positively charged nucleus.• The early nuclear model did not account

for all the atom’s mass• By 1920, Rutherford refined the concept of

the nucleus and concluded that the nucleus contained positively charged particles called protons.

(a) When a beam of alpha particles is directed at a thin gold foil, most particles pass through the foil undeflected, but a small number are deflected at large angles and a few bounce back toward the particle source. (b) A closeup view shows how most of an atom is empty space and only the alpha particles that strike a nucleus are deflected.

5) JAMES CHADWICK (1932)a) Discovered the

neutron

1) WHAT IS THE DIFFERENCE BETWEEN AN ATOM AND AN ELEMENT?a)Element = a substance that

cannot be broken down to other substances by a chemical reaction.

b)Atom: the smallest particle of an element that can exist either alone or in combination with other atoms.

2) 3 SUBATOMIC PARTICLES OF AN ATOM

2) SUBATOMIC PARTICLES OF AN ATOM

a)Subatomic particle = a particle smaller than an atom• Ex: proton, neutron, electron

2) SUBATOMIC PARTICLES OF AN ATOMb) An atom is composed of subatomic particles including protons, neutrons, and electrons (plus scientists have determined that protons and neutrons have their own structures and they are composed of quarks- these particles will not be covered since scientists do not yet understand if or how they affect chemical behavior).

c) Summary of subatomic particles

Particle Symbol

Location Relative

Electron

Charge

Relative Mass

Actual Mass (g)

Electron

e- In the space

surrounding

the nucleus

-1 1/1840 9.110 1028 g

Proton p In thenucleus

+1 1 1.673 1024 g

Neutron

n In thenucleus

0 1 1.675 1024 g

2) SUBATOMIC PARTICLES OF AN ATOM

d) Protons give the nucleus the positive charge

& determines the identity of an atome) Protons and neutrons have about the

same mass and are over 1840 times more massive than electrons.

2) SUBATOMIC PARTICLES OF AN ATOMf) Most of the mass of an atom is located in the

nucleus (protons & neutrons). (Nuclear forces hold the particles of a nucleus together. These forces only act over a very short range.)

g)Atoms are electrically neutral b/c of the number of protons EQUALS the number of electrons.

2) SUBATOMIC PARTICLES OF AN ATOMh)Nucleus = the positively charged, dense

central portion of an atoms• Contains nearly all of the atom’s mass,

but takes up a very small fraction of its volume

3) ATOMIC NUMBER & AVERAGE ATOMIC MASS

• Atomic number: the number of protons in the nucleiAlso represents the number of

electrons that an atom has since atoms are electrically neutral

Ex: Atomic number of Na is 11; Na has 11 protons and 11 electrons

19

K

Potassium

39.098

Atomic Number

Average Atomic Mass

• Average Atomic Mass: the weighted average of the atomic masses of naturally occurring isotopes of an element.

Isotopes = atoms of the same element with different number of NEUTRONS, therefore, different masses.

Periodic Table lists average atomic mass

3) ATOMIC NUMBER & AVERAGE ATOMIC MASS

Weighted averages account for the percentages of each isotope of a given element

Important b/c they indicate relative mass relationships in chemical reactions

3) ATOMIC NUMBER & AVERAGE ATOMIC MASS

• Ex: How average atomic mass is calculated

4) ISOTOPESa)Isotopes = atoms of the same element

with different number of neutrons, therefore, different masses.• Isotopes of an element all have the

SAME NUMBER of protons and electrons

• Named by their mass numbers

4) ISOTOPES• Mass number = the total number

of protons and neutrons in the nucleus of an isotope.

• Mass Number = # Protons + # Neutrons

4) ISOTOPES• Ex:

H-1

H-2

H-3

4) ISOTOPES

• Example 1: ALL carbon atoms have howmany protons? Most carbon atoms have 6 neutrons.

What is their mass number?12

6 (atomic number)

10) ISOTOPES• Some carbon atoms have 8 neutrons.

What is their mass number? 14

C-12 and C-14 are isotopes of carbon• Both isotopes have 6 electrons &

6 protons, but differ in the # of NEUTRONS!!

10) ISOTOPES• Example 2: How many neutrons

arein a sodium-23 atom?

• Example 3: How many protons, electrons and neutrons are there in an atom of chlorine-37?

17 protons; 17 electrons (atomic number is 17)

37 (protons + neutrons) – 17 protons = 20 neutrons

12

10) ISOTOPES• Hyphen Notation:

• Write the hyphen notation for a hydrogen isotope with a mass number of 3.

H-3 or Hydrogen-3• Nuclear Symbol Notation

• Write the nuclear symbol for a hydrogen isotope with a mass number of 3.

• Complete the following table.• Remember:

• atomic number = number of protons = number of electrons

• mass number = atomic number + number of neutrons

Nuclear Symbol

Atomic Number

Mass Number

Number of Protons

Number of Neutrons

Number of Electrons

3070

30 40 30

108

74

87

50

74 182 74

37 37 37

THE MOLE = SI unit for amount of substance

The Mole:→A mole simply represents a counting unit, much

in the same way that a dozen represents a set of twelve.

• Ex: Just as you can have a dozen cans of soda or a dozen donuts, you can have a mole of stars or a mole of water molecules.

• So a dozen eggs represents 12 eggs, a mole of carbon represents 6.022 X 1023

carbonatoms

• 6.022 X 1023 is also called “Avogadro’sNumber” = the number of particles

in exactly one mole of a pure substance.

The Mole:• There are 6.02 X 1023 atomic mass units (amu) in

one gram!!–This conversion allows us to use the

masses listed on the periodic table for both atomic mass and molar mass.

–Atomic mass = is the mass of one atom, is measured in atomic mass units (amu)

• Mass of C-12 is arbitrary assigned a mass of 12 amu

• 1/12 the mass of a carbon-12 atom is 1 amu

• Mass of any atom is expressed relative to the mass of one atom of carbon-12

The Mole:• Molar mass = is the mass of one mole of atoms

or molecules, is measured in grams.• Numerically equal to the atomic mass• 1 mole of a substance is 1 molar mass of

that substance• Ex: 1 mole of nitrogen is equal to 14.007g

of nitrogen!

Atomic & Molar Masses

Element & Symbol

Atomic Mass—Mass of 1 Atom

Molar Mass—Mass of 6.02 X 1023 atoms

Carbon (C) 12.011 amu 12.011 g

Helium (He) 4.0026 amu 4.0026 g

Copper (Cu) 63.546 amu 63.546 g

Potassium (K)

39.1098 amu 39.1098 g

The Mole:

• REMEMBER: one mole of atoms contains 6.02 X 1023 atoms. One mole of molecules contains 6.02 X 1023 molecules. One mole of formula units contains 6.02 X 1023 formula units. One mole of ions contains 6.02 X 1023 ions.

MOLES MASS (grams)• KEY IDEA: It is important to know

the following conversion:• The molar mass of an element is equal to the mass of 1 mole of atoms of that element.

• Example: 1 mole of zinc is equal to 65.39 grams of zinc!

• Use the factor label method (units will cancel out!)

MOLES MASS (grams)• Ex: What is the mass in grams of 3.5 mol

of the element copper (Cu)?

1. What is the mass in grams of

2.25 mol of iron (Fe)?

2. What is the mass in grams of

0.375 mol of K?

MASS (grams) MOLES

1 mol of an element = the molar mass of that element

MASS (grams) MOLES• Ex: A chemist produced 11.9g of Al.

How many moles of Al has been produced?

1. How many moles of Ca are contained in 5.0g of Ca?

2. How many moles of Au are contained in 3.6 x 10-10g of Au?

ATOMS MOLES

1 mol of an element = 6.02 x 1023

atoms of that element

ATOMS MOLES• Ex: How many moles of Ag are in 3.01 x

1023 atoms?

1. How many moles of Pb are equivalent to 1.5 x 1012

atoms?

2. How many moles of tin are equivalent to 2500 atoms?

3. How many atoms of Al are contained in 2.75 mol?

The Mole , Molar Mass, # of Atoms Practice Problems

• What is the conversion factor between moles and atoms?

1 mole = 6.02 X1023 atoms

The Mole , Molar Mass, # of Atoms Practice Problems

• What is the conversion factor between moles and mass?

1 mole = molar mass (grams)