Copyright © by Holt, Rinehart and Winston. All rights reserved.

ResourcesChapter menu

Table of Contents

Chapter 3 Atoms: The Building Blocks of Matter

Section 1 The Atom: From Philosophical Idea to Scientific Theory

Section 2 The Structure of the Atom

Section 3 Counting Atoms




Foundations of Atomic Theory

• Greeks- Democritus - atomos- “indivisible”

• Ideas based on opinion/thoughts

Chapter 3



Foundations of Atomic Theory

1790’s- quantitative analysis of chemical reactions

Several Laws that resulted from experimentation:

Law of Conservation of Mass

Law of Definite Proportions

Law of Multiple Proportions



Law of Conservation of Mass- mass is neither created nor destroyed in ordinary chemical reactions.

Section 1 The Atom: From Philosophical Idea to Scientific TheoryChapter 3



Law of Definite Proportions

A chemical compound always has the same mass ratio of elements

Example: SiO2 silicon dioxide (sand)

46.74% silicon 53.26 % oxygen

Mass of desired element X 100

Total mass of compound



Law of multiple proportions: The same two elements may combine in different whole number ratios- will give two different compounds

Section 1 The Atom: From Philosophical Idea to Scientific TheoryChapter 3




Dalton’s Atomic Theory 1808- based on experimental work- • 1. All matter is composed of extremely small particles

called atoms.• 2. Atoms of a given element are identical.• 3. Atoms cannot be subdivided, created, or

destroyed.• 4. Atoms of different elements combine in simple

whole-number ratios to form chemical compounds.• 5. In chemical reactions, atoms are combined,

separated, or rearranged.

Chapter 3




Modern Atomic Theory

• Not all aspects of Dalton’s atomic theory have proven to be correct. We now know that:

Chapter 3

• Atoms are divisible into even smaller particles• A given element can have atoms with different

masses.--- isotopes




The Structure of the Atom

• An atom is the smallest particle of an element that retains the chemical properties of that element.

Chapter 3




• Electrons- found by JJ Thomson (1897)– Used cathode ray tube– Negatively charged particles– Mass is small 1/1837 mass of hydrogen atom– Relative mass 0 amu (atomic mass unit)




Cathode ray tube

Chapter 3




Discovery of the Electron- J. J. Thomson

Cathode Rays and Electrons

Chapter 3

• Experiments in the late 1800s showed that cathode rays were composed of negatively charged particles.

• These particles were named electrons.

• “Plum-pudding” model- think chocolate chip cookie




• Protons- found by Ernest Rutherford(1918)– positively charged particles found in nucleus– Relative mass 1 amu (atomic mass unit)– # of protons- identifies element




• Neutrons- found by James Chadwick(1932)– neutral charge– particle found in nucleus– Relative mass 1 amu (atomic mass unit)



Properties of Subatomic Particles

Section 2 The Structure of the AtomChapter 3




Discovery of the Atomic Nucleus• More detail of the atom’s structure was provided in

1911 by Ernest Rutherford and his associates Hans Geiger and Ernest Marsden.

• The results of their gold foil experiment led to the understanding that the atom is mostly empty space with a small, dense, positive region.

• Rutherford called this positive bundle of matter the nucleus.

Chapter 3



Gold Foil Experiment




Gold Foil Experiment on the Atomic Level




Atomic Models- Visual Overview

Dalton

Thomson

Rutherford

Today?????



Determining the components of an atom

• Use the periodic table and given information in symbolic form to determine:

– Atomic number– Mass number– Numbers of protons, neutrons and electrons– Charge of an ion



Symbol you will see to represent an element

AXZ

A – mass number

Z – atomic number

X- symbol of the element




Atomic Number

• Atoms of the same element all have the same number of protons.

• The atomic number (Z) of an element is the number of protons of each atom of that element.

Chapter 3



Visual Concepts

Atomic Number

Chapter 3




Mass Number

• The mass number is the total number of protons and neutrons that make up the nucleus of an isotope.

Chapter 3



Visual Concepts

Mass Number

Chapter 3




Isotopes

• Isotopes are atoms of the same element that have different masses.

• The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons.

• Most of the elements consist of mixtures of isotopes.

Chapter 3




Designating Isotopes

• Hyphen notation: The mass number is written with a hyphen after the name of the element.

uranium-235

• Nuclear symbol: The superscript indicates the mass number and the subscript indicates the atomic number.

235 92 U

Chapter 3




Sample Problem A

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

Chapter 3




Additional Problems

Xe-135 ? p ? n ? e

25Mg +2 ? P ? N ? e

Chapter 3




Relative Atomic Masses

• The standard used 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.

Chapter 3




Average Atomic Masses of Elements

• Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.

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.

Chapter 3



Mass Spectrometer- http://www.chemguide.co.uk/analysis/masspec/masspec.GIF www.alevelchemistry.co.uk/Quizzes/images/mass_spectrometry




Average Atomic Masses of Elements, continuedCalculating Average Atomic Mass, continued

Chapter 3

• Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 601 amu, and 30.85% copper-65, which has an atomic mass of 64.927 794 amu.




Relating Mass to Numbers of Atoms

The Mole- “The Chemist’s Dozen”

Chapter 3

• Avogadro’s number 6.02 1023

– number of particles in exactly one mole of a pure substance.

Avogadro’s Number

• mole (abbreviated mol)• Enables small particles to be counted

• SI unit to describe quantity.




Relating Mass to Numbers of Atoms, continuedMolar Mass

Chapter 3

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

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

• The mass of one mole of a pure substance is called the molar mass of that substance.




Relating Mass to Numbers of Atoms, continued

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

Chapter 3

• For example, the molar mass of helium is

4.00 g He or 1 mol He

1 mol He 4.00 g He

Gram/Mole Conversions



Mole conversions

Molar mass in grams = 1 MOLE = 6.02 X1023 particles

Samples to try

moles and grams

moles and atoms

atoms and grams




Sample Problems- Moles to grams

What is the mass of one mole of aluminum?

What is the mass of 1 mol of chlorine?

Chapter 3




Sample Problem –

What is the mass of .455 moles of silver?

What is the mass of 2.50 X 10-4 moles of lead?

Chapter 3




Sample Problem- Moles to atoms

How many atoms are in 3.45 moles of copper?

Chapter 3



Moles to grams

• How much lead should I mass to get 1.67 moles?




Atoms to moles

How many moles of aluminum do you have with 4.75 1034 atoms of aluminum?

Chapter 3



Atoms to grams

• If I wanted 5.00 X 1015 atoms of gold, how many grams would I need?

Documents

Table of Contents