Unit 3Atomic Structure

Chemistry IMr. Patel

SWHS

Topic OutlineLearn Major IonsDefining the Atom (4.1)Subatomic Particles (4.2)Atomic Structure (4.2)Ions and Isotopes (4.3)Nuclear Chemistry (25.1)

Defining the AtomAtom – the smallest particle of an

element that retains its identityCan not see with naked eyeNanoscale (10-9 m)Seen with scanning

tunneling electronmicroscope

DemocritusDemocritus was a Greek to first

come up with idea of an atom.

His belief: atoms were indivisible and indestructible. = WRONG!

Atom comes from “atmos” - indivisible

Dalton’s Atomic Theory2000 yrs later, John Dalton used

scientific method to transform Democritus’s idea into a scientific theory

Dalton put his conclusions together into his Atomic Theory (4 parts)

Dalton’s Atomic Theory1. All elements are composed of tiny,

indivisible particles called atoms.

Dalton’s Atomic Theory2. Atoms of the same element are

identical. Atoms of different elements are different

Dalton’s Atomic Theory3. Atoms of different elements can

physically mix or chemically combine in whole number ratios.

Dalton’s Atomic Theory4. Chemical reactions occur when

atoms are separated, joined, or rearranged. Atoms of one element can never be changed into atoms of another element due to a chemical reaction.

The ElectronParticle with negative charge

Discovered by J.J. Thomson

Used cathode ray (electron) beam and a magnet/charged plate.

Millikan found the charge and mass

The Proton and NeutronAn atom is electrically neutral

If there is a negative particle then there must be positive particle

Proton – particle with positive charge

Chadwick discovered neutron – neutral charge

Thomson’s Atomic ModelElectrons distributed in a sea of

positive chargePlum Pudding Model

Rutherford’s Atomic ModelPerformed Gold-Foil ExperimentBeam of Alpha particles with positive

charge shot at thin piece of gold foilAlpha particles should have easily

passed through with slight deflection due to positive charge spread throughout.

Results: Most particles went straight through with no deflection. Some were deflected at large angles.

Rutherford’s Atomic ModelThe nucleus is the central part of the

atom containing protons and neutronsPositive chargeMost of the mass

Electrons are located outside the nucleusNegative chargeMost of the volume

Atomic Number An element is defined only by the

number of protons it contains

Atomic Number – number of protons

Number of protons = number of electronFor a neutral element

Identify the number of Protons1. Zinc (Zn)

2. Iron (Fe)

3. Carbon (C)

4. Uranium (U)

1. 30

2. 26

3. 6

4. 92

Mass NumberNucleus contains most of the massRounded Atomic MassMass Number – total protons and

neutrons

Number of neutron = Mass # – Atomic #

Identify # of Subatomic Particles

1. Lithium (MN = 7)

2. Nitrogen(MN = 14)

3. Fluorine(MN = 19)

**MN = Mass Number

1. 3 p+ , 3 e-, 4 n0

2. 7 p+ , 7 e-, 7 n0

3. 9 p+ , 9 e-, 10 n0

Differences in Particle NumberDifferent element: different number of

protons

Ions – same number of proton, different number of electrons

Isotope – same number of proton, different number of neutronsDifferent Mass Numbers

Two Notations for AtomsNuclear Notation

Write the element symbolOn left side, superscript = Mass NumberOn left side, subscript = Atomic Number

Isotope –Hyphen NotationWrite full name of elementOn right side, put a dashOn right side put Mass Number after dash

Hydrogen - 3

Ex: Three isotopes of oxygen are oxygen-16, oxygen-17, and oxygen-18.

Write the nuclear symbol for each.

Ex: Three isotopes of chromium are chromium-50, chromium-52, and chromium-53. How many neutrons are in each isotope?

Ex: Calculate the number of neutrons for 99

42Mo.

Ex: Calculate the number of neutrons for 238

92U.

Ex: Classify the following atoms21

45X 2345X 20

45X.

Ex: Classify the following atoms196

79X 19580X 195

78X.

Atomic MassAtomic Mass Unit (amu) – one-twelfth

of the mass of the carbon-12 atom

Different isotopes have different amu (mass) and abundance (percentage of total)

Atomic Mass – weighted average mass of the naturally occurring atoms.Isotope MassIsotope Abundance

Atomic MassPercent Abundance – the number of desired

particles in 100 total particles of sampleAllows for comparison to any sample set

Relative Abundance – the number of desired particles in the sample usedSpecific to the sample used; not useful in

comparisonConvert % abundance to a decimal =

relative abundance

Desired particlesTotal particles in sample

% Ab = x 100%

Atomic MassBecause abundance is considered, the

most abundant isotope is typically the one with a mass number closest to the atomic mass.

Example, Boron occurs as Boron-10 and Boron-11. Periodic Table tells us Born has atomic mass of 10.81 amu.Boron-11 must be more

abundant

Calculating Atomic MassConvert the Percent Abundance to

Relative Abundance (divide by 100)

Multiple atomic mass of each isotope by its relative abundance

Add the product (from step above) of each isotope to get overall atomic mass.

Ex: If there are 100 black beans, 27 pinto beans, and 173 lima beans in the container, what is the percent

abundance of the container by bean? Relative abundance?

Ex: Calculate the atomic mass for bromine. The two isotopes of bromine have atomic masses and percent abundances of 72.92 amu (50.69%) and 80.92 amu

(49.31%).

Ex: Calculate the atomic mass for X. The four isotopes of X have atomic masses and percent abundances of 204 amu (1.4%), 206 amu (24.1%), 207 amu (22.1%), and

208 amu (52.4%).

Ex: Calculate the atomic mass for H. The three isotopes of H have atomic masses and percent

abundances of 27 amu (85%), 26 amu (10%), and 28 amu (5%).

Nuclear Radiation

Radioactivity – nucleus emits particles and rays (radiation)

Radioisotope – a nucleus that undergoes radioactive decay to become more stable

An unstable nucleus releases energy through radioactive decay.

Nuclear Radiation

Nuclear force – the force that holds nuclear particles together Very strong at close distances

Of all nuclei known, only a fraction are stableDepends on proton to neutron ratioThis region of stable nuclei called band

of stability

Half Life

Half Life – the time required for one-half the sample to decayCan be very short

or very long

Symbol Element Radiation Half-LifeDecay

Product

U-238Uranium-

238alpha

4,460,000,000 years

Th-234

Th-234Thorium-

234beta 24.1 days Pa-234

Pa-234Protactiniu

m-234beta

1.17 minutes

U-234

U-234Uranium-

234alpha

247,000 years

Th-230

Th-230Thorium-

230alpha

80,000 years

Ra-226

Ra-226Radium-

226alpha

1,602 years

Rn-222

Rn-222 Radon-222 alpha 3.82 days Po-218

Po-218Polonium-

218alpha

3.05 minutes

Pb-214

Pb-214 Lead-214 beta 27 minutes Bi-214

Bi-214Bismuth-

214beta

19.7 minutes

Po-214

Po-214Polonium-

214alpha

1 microseco

ndPb-210

Pb-210 Lead-210 beta 22.3 years Bi-210

Bi-210Bismuth-

210beta 5.01 days Po-210

Po-210Polonium-

210alpha 138.4 days Pb-206

Pb-206 Lead-206 none stable (none)

Ex: The original amount of sample was 100 g. The amount currently remaining is

25 g. How many half-lives has gone by?

Ex: The original amount of sample was 100 g. The amount currently remaining is 25 g after 30 minutes. What is the half life?

Ex: The original amount of sample was 100 g. The amount currently remaining is 6.25 g. The half life

is 50 years. How much time has passed?

Nuclear ReactionsDeals with nucleus

Can end up with new atoms/elements

Mass is not strictly conserved Mass DefectE = mc2

Deals with electrons

Atoms/elements remain unchanged – rearranged

Mass is strictly conserved

Nuclear vs. Chemical Reactions

Chemical Reactions

Types of Radiation

Alpha Radiation (Helium Atom)Low penetrating powerPaper shielding

Beta Radiation (Electron)Moderate penetrating powerMetal foil shielding

Gamma Radiation (Pure energy)Very high penetrating powerLead/concrete shielding

Nuclear Decay Equations

Transmutation – conversion from one element to another through a nuclear reactionOnly occur by radioactive decayOnly when nucleus bombarded with a particle

Emissions – given offAlpha Emission, Beta Emission, Positron EmissionPositron = beta particle with a positive charge

Captures – taken inElectron Capture

Ex: Show a Beta Emission of Copper-66.

Ex: Show an Electron Capture of Nickel-59.

Ex: Show a Positron Emission of Boron-8.

Ex: Show an Alpha Emission of Thorium-232.