Unit 3: “Atomic Structure”

Chemistry: Mr. Blake/Mr. Chemistry: Mr. Blake/Mr. GowerGower

I. Atomic Structure

NOTE: The Greek philosopher __________ (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”)

– He believed that atoms were ________ and _____________

– His ideas did agree with later scientific theory, but did not explain chemical behavior, and was ________________

_______________– but just philosophy

Democritus

indivisibleindestructible

not based on the scientific method

A. Section 4.1 Defining the Atom

1. Dalton’s Atomic Theory (experiment based!)

c. Atoms of different elements combine in simple whole-number ratios to form chemical _________d. In chemical reactions, atoms are combined, separated, or rearranged – but ____ changed into atoms of another element.

a. All elements are composed of tiny indivisible particles called _____.b. Atoms of the same element are _______. Atoms of any one element are _______ from those of any other element. John Dalton

(1766 – 1844)

atomsidentical differen

t

compounds

never

Problems with Dalton’s Atomic Theory?1. matter is composed of indivisible particles

Atoms Can Be Divided, but only in a nuclear reaction2. all atoms of a particular element are identical

Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)! Different elements have different atoms. YES!

3. atoms combine in certain whole-number ratiosYES! Called the Law of Definite Proportions

4. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements.Yes, except for nuclear reactions that can change atoms of one element to a different element

2. Sizing up the Atoma. Elements are able to be subdivided into smaller and smaller particles – these are the _____, and they still have __________ of that elementb. If you could line up 100,000,000 copper atoms in a single file, they would be approximately ________c. Despite their ________, individual atoms ___ observable with instruments such as scanning tunneling (electron) microscopes

atomsproperties

1 cm longsmall

sizeare

B. Section 4.2 Structure of the Nuclear Atom

NOTE: One change to Dalton’s atomic theory is that _______________ into subatomic particles:NOTE: ________________________ are examples of these fundamental particlesNOTE: There are many other types of particles, but we will study these three

atoms are divisible

Electrons, protons, and neutrons

1. Discovery of the Electrona. In 1897, J.J. Thomson used a _______ _______ to deduce the presence of a negatively charged particle: the _______

cathoderay tube

electron

Modern Cathode Ray Tubes

Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Television Computer Monitor

2. Mass of the Electron

a. 1916 – _____________ determines the _____ of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge

The oil drop apparatus

Mass of the electron is 9.11 x 10-28 g

Robert Millikanmass

3. Conclusions from the Study of the Electron:

a. Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.b. Atoms are neutral, so there must be ______________ in the atom to balance the negative charge of the electronsc. _______________________ that atoms must contain other particles that account for most of the mass

positive particles

Electron have so little mass

d. ____________ in 1886 observed what is now called the “______” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron)e. 1932 – _____________ confirmed the existence of the “______” – a particle with ________, but a mass nearly ____ to a proton

Eugen Goldstein proton

James Chadwick neutron

no chargeequal

4. Subatomic Particles

Particle

Charge

Mass (g) Location

Electron (e-)

-1 9.11 x 10-28 Electron cloud

Proton (p+) +1 1.67 x 10-24 Nucleus

Neutron (no)

0 1.67 x 10-24 Nucleus

NucleusNucleus

Electron Electron cloudcloud

5. Thomson’s Atomic Model

a. Thomson believed that the ________ were like plums embedded in a positively charged “pudding,” thus it was called the “__________” model.

J. J. Thomson

electrons

plum pudding

6. Ernest Rutherford’sGold Foil Experiment - 1911

a. Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foilb. Particles that hit on the detecting screen (film) are recorded

7. Rutherford’s problem:a. In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?

Target #1

Target #2

b. The Answers:

Target #1 Target #2

8. Rutherford’s Findings

1) The nucleus is _____2) The nucleus is _____3) The nucleus is _______ charged

a. Most of the particles passed right through

b.A few particles were deflectedc. VERY FEW were greatly deflected

“Like howitzer shells bouncing off of tissue paper!”

d. Conclusions:smalldensepositively

9. The Rutherford Atomic Model

a. Based on his experimental evidence:1) The atom is mostly empty space2) All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “______”

3) The nucleus is composed of ______ and ________ (they make the nucleus!)4) The electrons distributed around the nucleus, and occupy most of the ______5) His model was called a “___________”

nucleus protons

neutrons

volumenuclear

model

1. Atomic Numbera. Atoms are composed of _______ protons, neutrons, and electronsb. How then are atoms of one element different from another element?c. Elements are different because they contain different numbers of ________d. The “____________” of an element is the _______________ in the nucleuse. ___________________________

identical

PROTONSatomic numbernumber of protons

# protons in an atom = # electrons

C. Section 4.3: Distinguishing Among Atoms

2. Definition: Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.

Element # of protons Atomic # (Z)

Carbon

Phosphorus

Gold

6 6

15

15

79

79

3. Mass NumberDefinition: Mass number is the number of protons and neutrons in the nucleus of an isotope:

Mass # = p+ + n0

Nuclide p+ n0 e- Mass #

Oxygen - 10

- 33 42

- 31 15

8 8 1818

Arsenic 75 33 75

Phosphorus 15 3116

4. Nuclear/Complete Symbols

a. Contain the symbol of the element, the mass number and the atomic number.

X Massnumber

Atomicnumber

Subscript →

Superscript →

a. Find each of these: a. Find each of these:

1) number of protons1) number of protons

2) number of 2) number of neutronsneutrons

3) number of 3) number of electronselectrons

4) Atomic number4) Atomic number

5) Mass Number5) Mass Number

Br80 35

5. Symbols

35

453535

80

b. If an element has an b. If an element has an atomic number of 34 and a atomic number of 34 and a mass number of 78, what is mass number of 78, what is the: the:

1)1) number of protonsnumber of protons

2)2) number of neutronsnumber of neutrons

3)3) number of electronsnumber of electrons

4)4) complete symbolcomplete symbol

34

4434

Se78 34

d. If an element has 78 d. If an element has 78 electrons and 117 electrons and 117 neutrons what is the neutrons what is the

1)1) Atomic numberAtomic number

2)2) Mass numberMass number

3) number of protons3) number of protons

4) complete symbol4) complete symbol

78

195

78

Pt195 78

a. Dalton was wrong about all elements of the same type being identicalb. Atoms of the same element can have different numbers of _______.c. Thus, different mass numbers.d. These are called _______.

neutrons

isotopes

6. Isotopes

Isotopes

e. _____________(1877-1956) proposed the idea of isotopes in 1912f. _______ are atoms of the ____ ______ having different masses, due to varying numbers of neutrons.g. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials.

Frederick Soddy

Isotopes sameelement

a. We can also put the mass number after the name of the element:b. Examples:

•carbon-12•carbon-14•uranium-235

7. Naming Isotopes

c. _______ are atoms of the ___________ having ________ masses, due to varying numbers of neutrons.

Isotope Protons

Electrons

Neutrons

Nucleus

Hydrogen–1

(protium)

1 1 0

Hydrogen-2

(deuterium)

1 1 1

Hydrogen-3

(tritium)

1 1 2

Isotopes same elementdifferent

8. Isotopesa. Elements occur in nature as _______ of _______.

b. Isotopes are atoms of the same element that differ in the _______ _______.

mixturesisotopes

number ofneutrons

a. ____ are atoms or groups of atoms with a positive or negative

charge.

b. _________ an electron from an atom gives a _____ with a

____________

c. ______ an electron to an atom gives an _____ with a ____________.

d. To tell the difference between an atom and an ion, look to see if

there is a charge in the _________! Examples: Na+ Ca+2 I- O-2

Na Ca I O

Ions

Taking away

cation positive charge

Adding anion

negative charge

superscript

9. IONSIONS

e. A cation forms when an atom loses one or more electrons.

f. An anion forms when an atom gains one or more electrons

Mg --> Mg2+ + 2 e- F + e- --> F-

metalsmetals (Mg) (Mg) lose electrons lose electrons ---> ---> cationscations

nonmetalsnonmetals (F) (F) gain electronsgain electrons ---> ---> anions

NOTE: In General……

Learning Check – Counting

State the number of protons, neutrons, and electrons in each of these ions.

39 K+ 16O -2 41Ca +2

19 8 20

#p+ ______ ______ _______

#no ______ ______ _______

#e- ______ ______ _______

19

20

18

8

8

10

20

21

18

One Last Learning Check

Write the nuclear symbol form for the following atoms or ions:

A. 8 p+, 8 n, 8 e- ___________

B. 17p+, 20n, 17e- ___________

C. 47p+, 60 n, 46 e- ___________

O 168

Cl 3717

Ag 10747

Charges on Common Charges on Common IonsIons

Charges on Common Charges on Common IonsIons

-1-2-3+1

+2

By losing or gaining e-, atom has same By losing or gaining e-, atom has same number of e-number of e-’’s as nearest Group 8A atom.s as nearest Group 8A atom.

Example: A student has a test percentage of 78%; a lab percentage of 92%; and has completed homework at 100%.

Her weighted average grade is computed as(78% X 0.6) + (92% X 0.20) + (100% X 0.20)= 84.5

a. How heavy is an atom of oxygen? It depends, because there are

different _____ of oxygen atoms.b. We are more concerned with the _________________.c. This is based on the abundance (percentage) of each variety of that element in nature.d. We don’t use grams for this mass because the numbers would be too small.

kinds

average atomic mass

10. Atomic Mass

a. Instead of grams, the unit we use is the ______________ (amu)b. It is defined as one-twelfth the mass of a carbon-12 atom.c. Carbon-12 chosen because of its _____ ______. d. Each isotope has its own atomic mass, thus we determine the average from percent abundance.

Atomic Mass Unit

isotopepurity

11. Measuring Atomic Mass

a. Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results.b. If not told otherwise, the mass of the isotope is expressed in _____________ (amu) atomic mass units

12. To calculate the average:

Isotope Symbol Composition of the nucleus

% in nature

Carbon-12

12C 6 protons6 neutrons

98.89%

Carbon-13

13C 6 protons7 neutrons

1.11%

Carbon-14

14C 6 protons8 neutrons

<0.01%

Atomic mass is the average of all the naturally occurring isotopes of that element.

Carbon = 12.01

13. Atomic Masses

- Page 117

Question

Solution

Answer

Knowns and Unknown

14. Calculate the atomic mass of carbon.

a. Isotopes % abundance Atomic mass

Carbon-12 98.89% 12.000 amu

Carbon-13 1.11% 13.003 amu

.b. Lithium has two isotopes. If lithium-6 has a mass of

6.015 and 7.42 % occurrence, what is the % abundance and mass of lithium -7?

Atomic mass = (%) (mass) + (%) (mass) + …..

Atomic mass = (0.9889) (12.000) + (0.0111) (13.003)

= 11.87 + .144 = 12.01 amu

6.941 = (0.0742)(6.015) + (0.9258)(x)

= 7.015 amu

a. A “periodic table” is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties1) The periodic table allows you to easily compare the properties of one element to another

15. The Periodic Table: A Preview

b. Each horizontal row (there are 7 of them) is called a _____c. Each vertical column is called a ____________1) Elements in a _____ have similar chemical and physical properties2) Identified with a number and either an “A ” or “B”

period

group or family group

• A few elements, such as gold and copper, have been known for thousands of years - since ancient times

• Yet, only about __ had been identified by the year 1700.

• As more were discovered, chemists realized they needed a way to ________ the elements.

13

organize

II. The Periodic TableA. Section 6.1: Organizing the ElementsNOTES:

•Chemists used the _________ of elements to sort them into groups.

• In 1829 J. W. Dobereiner arranged elements into _____ – groups of three elements with similar properties

•One element in each triad had properties intermediate of the other two elements

properties

triads

a. By the mid-1800s, about 70 elements were known to existb. Dmitri _________ – a Russian chemist and teacherc. Arranged elements in order of _________________d. Thus, the first “Periodic Table”

Mendeleev

increasing atomic mass

1. Mendeleev’s Periodic Table

a. ___________ for yet undiscovered elementsb. When they were discovered, he had made good predictionsc. But, there were problems:

–Such as Co and Ni; Ar and K; Te and I

He left blanks

2. Mendeleev

a. In 1913, Henry ______ – British physicist, arranged elements according to increasing ____________b. The arrangement used todayc. The symbol, atomic number & mass are basic items included-textbook page 162 and 163

Moseley

atomic number

3. A Better Arrangement

a. When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.b. Horizontal rows = ______1) There are __ periodsc. Vertical column = _____ (or family)1) Similar physical & chemical prop.2) Identified by number & letter (IA, IIA)

periods7

group

4. The Periodic Law

5. Areas of the Periodic TableThree classes of elements are:

1) _____, 2) ________, and 3) _________

1) Metals: _______ conductors, have luster, ductile, malleable

2) Nonmetals: generally brittle and non-lustrous, poor conductors of ____ and electricity

metals nonmetalsmetalloid

s electrical

heat

Areas of the Periodic Table

• Some nonmetals are _____ (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br)

• Notice the heavy, stair-step line?3) _________: border the line-2 sides

– Properties are __________ between metals and nonmetals

gases

Metalloids

intermediate

Squares in the Periodic Table• The periodic table displays the ______

and _____ of the elements, along with information about the structure of their atoms:•Atomic ______ and atomic _____•Black symbol = solid; red = gas; ____ _____ (from the Periodic Table on our classroom wall)

symbolsnames

number

mass blue

=liquid

Groups of Elements - Family Names• Group IA (1) – __________

– Forms a “base” (or alkali) when _______ with water (not just dissolved!)

• Group 2A (2)– ________________– Also form bases with water; do not dissolve well,

hence “earth metals”

• Group 7A (17) – _______– Means “salt-forming”

• Group 8A (18) – _________– Nonreactive because of their electron configuration

alkali metals

alkaline earth metals

halogensnoble gases

reacting

ELEMENTS THAT EXIST AS ELEMENTS THAT EXIST AS DIATOMICDIATOMIC MOLECULES MOLECULES

ELEMENTS THAT EXIST AS ELEMENTS THAT EXIST AS DIATOMICDIATOMIC MOLECULES MOLECULES

Remember:

HOFBrINClThese

elements only exist as PAIRS. Note that when

they combine to make

compounds, they are no

longer elements so they are no

longer in pairs!

Nuclear ChemistryNuclear Chemistry

Chemistry – Unit 4Chemistry – Unit 4Chapter 25Chapter 25

Mass Defect

• Difference between the mass of an atom and the mass of its individual particles.

4.00260 amu 4.03298 amu

Nuclear Binding Energy

• Energy released when a nucleus is formed from nucleons.

• High binding energy = stable nucleus.

E = mc2

E: energy (J)m:mass defect (kg)c: speed of light

(3.00×108 m/s)

Nuclear Binding Energy

Unstable nuclides are radioactive and undergo radioactive decay.

He42

Types of Radiation

• Alpha particle ()– helium nucleus paper2+

• Beta particle (-)– electron e0

-1 1-lead

• Positron (+)– positron e0

1 1+

• Gamma ()– high-energy photon 0

concrete

Nuclear Decay

• Alpha Emission

He Th U 42

23490

23892

parentnuclide

daughternuclide

alphaparticle

Numbers must balance!!

Nuclear Decay

• Beta Emission

e Xe I 0-1

13154

13153

electron• Positron Emission

e Ar K 01

3818

3819

positron

Nuclear Decay

• Electron Capture

Pd e Ag 10646

0-1

10647

electron• Gamma Emission

– Usually follows other types of decay.

• Transmutation – One element becomes another.

IQ# 1

1.Balance the following equations:

HeNp 42

23793

ePo1

021284 Bi

21283

Pa23391

Nuclear Decay

• Why nuclides decay…– need stable ratio of neutrons to protons

He Th U 42

23490

23892

e Xe I 0-1

13154

13153

e Ar K 01

3818

3819

Pd e Ag 10646

0-1

10647

Belt of Stability Belt of Stability and Radioactive and Radioactive DecayDecay

Half-life

• Half-life (t½)– Time required for half the atoms of a

radioactive nuclide to decay.– Shorter half-life = less stable.

Half-life

nif mm )( 2

1

mf: final massmi: initial massn: # of half-lives

Half-life• Fluorine-21 has a half-life of 5.0 seconds.

If you start with 25 g of fluorine-21, how many grams would remain after 60.0 s?

GIVEN:

t½ = 5.0 s

mi = 25 g

mf = ?

total time = 60.0 sn = 60.0s ÷ 5.0s =12

WORK:mf = mi (½)n

mf = (25 g)(0.5)12

mf = 0.0061 g

Example: How much of a 500. g sample of Uranium-235 would be left after five half-lives?

(n = # of half-lives)Mi = 500 gn = 5Mf = ?

mf = mi (½)n

mf = (500 g)(0.5)5

mf = 15.6 g

Example: A 16.00 mg sample of Radon-222 decays to 0.250 mg after 24 hours. Determine the half-life.

h 4.0 h 4 6

h 24

16→ 8 → 4 → 2 → 1 → 0.5 → 0.250 = 6 half lives

Example: The half-life of molybdenum-99 is 67 hours. How much of a 1.000 mg sample is left after 335 hours?

Mi = 1.000 mgHalf-life = 67 h

Rxn time = 335 hMf = ?

n = 335 / 67 = 5

mf = mi (½)n

mf = (1.000 mg)(0.5)5

mf = 0.03125 mg

Learning Check!

The half life of I-123 is 13 hr. How much of a 64 mg sample of I-123 is left after 39 hours?

Mi = 64 mgn = 3

Mf = ?

mf = mi (½)n

mf = (64 mg)(0.5)3

mf = 8.0 mg

Half-life LabProcedure:1. Each lab group will acquire a sample of 50 pennies in

a cup.2. Count pennies to make sure you have 50 pennies.3. Enter “50” in Shake # 0 row for Trial 1, 2, & 3

and “150” for (Sum of) of trials.4. 4. Shake the cup of pennies. Pour the pennies on to

the lab bench.5. Remove all pennies that land on “heads”. They have

decayed.6. Count only the remaining pennies (the pennies that

landed on “tails”). Record data.7. Place only the remaining pennies (“tails”) into the

cup and shake again. Repeat steps 4-7 until all pennies have decayed.

8. Repeat the process two more times and record data under Trial 2 & 3.

• Data: Collect data for three trials in the table.

• Data Analysis: Prepare a graph to represent the decay of your sample ( of trials (y-axis) vs. Shake # (x-axis))

• Prepare a graph in your lab book: Graph the # of undecayed atoms ( of trials) (y-axis) versus the Shake # (x-axis). Label the x and y axes, including units (if applicable). Make graph large (at least 2/3 pg.). Draw a best fit curve that represents your data. Use a Ruler!

• Plot the Shake # for the of trials using the best fit curve.

• Determine the “half life” of your sample in terms of # of shakes using your graph.

Graphing the ResultsImportant !!

• Title every graph and label each axis (include units)

• Graphs should be at least 2/3 page• Use a ruler• Circle all data points• Use a best-fit line (no “connect the dots”!)• Find the average half-life (in # of trials) of

your sample by interpolating your curve at exactly 75, 37.5, and 18.75 pennies undecayed)

Half-Life LabΣ

of

Tri

als

Shake #

·

·

·

·

Use Ruler for axisLabel Axis

Best Fit CurveConvenient #’s

Circle Data Points

At least 2/3 of pgTitle

75

37.5

18.75

150

01 2 3 4 5 6 7 8 9 100

1 + 1 + 1.2 = 3.2

3~ 1.1 shake

Half-life = 1.1 shake

1 2 3

F ission

• splitting a nucleus into two or more smaller nuclei

• 1 g of 235U = 3 tons of coal

U23592

F ission• chain reaction - self-propagating reaction• critical mass -

mass required to sustain a chain reaction

Fusion• combining of two nuclei to form one nucleus of larger mass• thermonuclear reaction – requires temp of 40,000,000 K to sustain• 1 g of fusion fuel =

20 tons of coal• occurs naturally in

stars

HH 31

21

Fission vs. Fusion

• 235U is limited• danger of

meltdown• toxic waste• thermal pollution

• fuel is abundant• no danger of

meltdown• no toxic waste• not yet sustainable

FISSION

FUSION

Nuclear Power

• Fission ReactorsCooling Tower

Nuclear Power

• Fission Reactors

Nuclear Power

• Fusion Reactors (not yet sustainable)

Nuclear Power

• Fusion Reactors (not yet sustainable)

Tokamak Fusion Test Reactor

Princeton University

National Spherical Torus

Experiment

Synthetic Elements• Transuranium Elements

– elements with atomic #s above 92– synthetically produced in nuclear reactors and accelerators– most decay very rapidly

Pu He U 24294

42

23892

Radioactive Dating

• half-life measurements of radioactive elements are used to determine the age of an object

• decay rate indicates amount of radioactive material

• EX: 14C - up to 40,000 years238U and 40K - over 300,000

years

Nuclear Medicine

• Radioisotope Tracers– absorbed by specific organs and used to

diagnose diseases

• Radiation Treatment– larger doses are used

to kill cancerous cells in targeted organs

– internal or external radiation source

Radiation treatment using

-rays from cobalt-60.

Nuclear Weapons

• Atomic Bomb– chemical explosion is used to form a

critical mass of 235U or 239Pu– fission develops into an uncontrolled

chain reaction

• Hydrogen Bomb– chemical explosion fission fusion– fusion increases the fission rate– more powerful than the atomic bomb

Others

• Food Irradiation radiation is used to kill bacteria

• Radioactive Tracers– explore chemical pathways– trace water flow– study plant growth, photosynthesis

• Consumer Products– ionizing smoke detectors - 241Am