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Chapter Overview

The Structure The Structure of the Atomof the Atom

Chapter 4

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4.1 - Early Theories of the Atom

4.2 - Subatomic Particles

4.3 - How Atoms Differ

4.4 - Unstable Nuclei & Radioactivity

Ann Dau – CHEMISTRY 112

4.1 Early Theories of Matter• Dalton’s Atomic Theory

1. All matter is composed of extremely small particles called atoms.

2. All atoms of a given element are identical. Atoms of different elements are different from one another.

3. Atoms cannot be created or divided into smaller particles or destroyed.

4. Different atoms combine in simple whole number ratios to form compounds.

5. In a chemical reaction, atoms are separated, combined, or rearranged.

Dalton’s Atomic ModelAtom - the smallest

particle of an element that retains the

properties of the element.

4.1 Early Theories of Matter

• E. Goldstein discovered the proton in 1886. • J.J. Thomson discovered the electron in

1897 during cathode ray tube experiments in the late 1890’s.

4.1 Early Theories of Matter

• Robert A. Millikan determined the mass and charge of the electron in 1916.– one unit of negative charge– mass is 1/1840 of a hydrogen atom

4.1 Early Theories of Matter

• In 1911 Ernest Rutherford discovered the nucleus during his gold foil experiment.

4.1 Early Theories of Matter

Ernest Rutherford’s gold foil experiment.

4.1 Early Theories of Matter

• Neils Bohr developed the planetary model of the atom– Electrons are in a particular

path have a fixed energy

– Energy level-region around a nucleus where the electron is likely to be moving

4.1 Early Theories of Matter

• Erwin Schrodinger developed the Quantum Mechanical Model– Describes the electronic

structure of the atom as the probability of finding electrons within certain regions of space

4.1 Early Theories of Matter

• James Chadwick discovered the neutron in 1932.

• In 1913 Henry Mosley used X-rays to count the number of protons in the atomic nuclei of different atoms.

The Atomic Theory

All elemental substances are made up of tiny indivisible particles called atoms.

Atoms of different elements have different properties.

In an ordinary chemical reaction, atoms of elements maintain their identities.

Compounds are made when atoms of different elements combine.

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Parts of the Atom

Protons give atoms their identity!

Atomic Structure

NucleusProtons and neutrons

Both about the same mass (Table 2.1 Text)

About 10-13 cm radius for gold atom

Most of atomic mass found in nucleus

Outer peripheryelectrons

Less massive

Electronic structure covered later

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4.2 Subatomic Particles & the Nuclear Atom

• Located within the Nucleus– Proton (p+)

• Positively charged particle (each carries a charge of +1)• Relative mass = 1 amu• Actual mass = 1.673 X 10-27 kg

– Neutron (n0)• Neutrally charged particle• Relative mass = 1 amu• Actual mass = 1.675 X 10-27 kg• Serves as the glue that holds the nucleus together as well as a buffer

between the charges of protons and electrons

Subatomic Particles

• Located outside the nucleus in the electron cloud– Electron

• Negatively charged particle (each carries a charge of -1)• Relative mass = 1/1840 amu• Actual mass = 9.11 X 10-31 kg• The electron is the part of the atom that will function in

bonding and reactions

4.3 How Atoms Differ

• Atomic Number– the number of protons in the nucleus of an atom– indicated at the top of the element’s block on the

periodic table 88

OO15.99915.999

1212

MgMg24.30524.305

Oxygen has an atomic number of 8

There are 8 protons in an atom of Oxygen

Magnesium has an atomic number of 12

There are 12 protons in an atom of Magnesium

Isotopes

• Atoms of the same element with the same number of protons, but different numbers of neutrons

• Since the atoms have different numbers of neutrons, they also have different mass numbers– Mass number = # of protons + # of neutrons

Abbreviating Isotopes

• Hyphen Notation– Simply write the Name of the atom, put a

hyphen, and then write the mass number• Carbon-12 vs. Carbon-14

– Carbon 12 has 6 protons and 6 neutrons– Carbon 14 has 6 protons and 8 neutrons

• Nuclear Designation– Element symbol is written in the center– Mass number goes in the upper left corner– Atomic number goes in the lower left corner C

12

6

Different Isotopes• Identify the number of protons, neutrons, and

electrons each of the following have.

Boron-10

Boron 11

35Cl 66Zn17 30

35

Br79.904

p+: ________ no: ________ e-: ________

p+: ________ no: ________ e-: ________

p+: ________ no: ________ e-: ________

p+: ________

no: ________

e-: ________

p+: ________

no: ________

e-: ________

Sample Problem 2.2 Determining the Number of Subatomic Particles in the Isotopes of an Element

PROBLEM: Silicon(Si) is essential to the computer industry as a major component of semiconductor chips. It has three naturally occurring isotopes: 28Si, 29Si, and 30Si. Determine the number of protons, neutrons, and electrons in each silicon isotope.

Calculating Atomic Mass

• Mass Number– the number of protons + neutrons in a given isotope

• Atomic Mass– The weighted average mass of all of the isotopes of

that element[(Mass of isotope A)(percent abundance )] + [(Mass of isotope B)(percent abundance)]

Practice Calculating Atomic Mass

Calculate the atomic mass of helium given the following information:

There are two naturally occurring isotopes of helium:

Isotope % Abundance Mass

helium-3 0.0001 3.0160

helium-4 99.9999 4.0026

Practice Calculating Atomic Mass

There are two naturally occurring isotopes of helium:Isotope % Abundance Mass

helium-3 0.0001 3.0160 helium-4 99.9999 4.0026

(3.0160 x 0.000 001) + (4.0026 x 0.999999) 0.000 003 0160 + 4.002595997 0.000 003 0160 + 4.0026

= 4.002603016 = 4.0026

Practice Calculating Atomic Mass

There are three naturally existing isotopes of silicon: silicon-28, silicon-29, and silicon-30. Their percents of natural abundance is listed respectfully: 92.21 %, 4.70 %, and 3.09 %.

Calculate the average atomic mass of silicon and express your answer in 4 significant digits.

Sample Problem 2.3 Calculating the Atomic Mass of an Element

Silver (Ag: Z = 47) has 46 known isotopes, but only 2 occur naturally, 107Ag and 109Ag. Given the following data, calculate the atomic mass of Ag:

Isotope Mass(amu) Abundance(%)

107Ag109Ag

106.90509

108.90476

51.84

48.16

Atoms, Ions, and IsotopesIons and isotopes are atoms…they are

specific types of atoms.

Ions: atoms that have a charge; either positive or negative

What would allow an atom to be an ion?

Isotopes: a relationship between two atoms in which they have the same number of protons and different numbers of neutrons.

Are isotopes technically atoms of the same element?

Ions

• Atoms that have gained electrons to become___________________ charged.

• Atoms that have lost electrons to become___________________ charged.

Key: Has the identity of the atoms changed?Why?

4.4 Unstable Nuclei & Radioactive Decay

• Nuclear Reactions– reactions that involve a change in the nucleus of an atom.

• Radioactivity– the spontaneous release of radiation.

• Radiation– rays and particles emitted by radioactive materials

• Radioactive atoms emit radiation because their nuclei are unstable.

• There are three main types of radiation– Alpha decay– Beta decay– Gamma decay

Alpha (α) radiation

• two protons and 2 neutrons

• Positive charge

• Symbols: 4 4He

2

2

• reduces the atomic number by 2

• reduces the mass by 4

Beta (β) radiation

• Fast moving electron

• Negative charge

• Symbols: 0 -1

• increases the atomic number by 1

• does not change the mass

Gamma (γ) radiation

• high energy radiation

• released with alpha and beta radiation

• symbol: 0 0

• does not change the mass or atomic number

Half lives

• The time it takes for 1/2 of the mass of the isotope to be decayed.

• If I have a 60g sample and the half life is 2 years, how long will it take for there to be 7.5g left of the sample?

60g 30g 15g 7.5g 2 years 4 years 6 years

So, it takes 6 years for the 60g sample to decay into 7.5g.So, it takes 6 years for the 60g sample to decay into 7.5g.