Chapter 3 Atoms: The Building Blocks of Matter

Lesson 1.1: Early Atomic Theory Learning Target: I will understand the history and structure of the atom. Success Criteria: I can describe the experiments and discoveries of Democritus, Dalton, Thomson, and Rutherford. I can describe Thomson’s and Rutherford’s model of the atom and the know the subatomic particles involved.

Democritus (400 B.C.E) Greek philosopher Hypothesized: Nature has a basic indivisible particle of which everything is made of Called this particle an atom Greek “atomos” = indivisible

John Dalton ( 1808 A.D.) English schoolteacher – liked nature and weather Dalton’s Atomic Theory: 1. All matter is composed of extremely small particles called atoms 2. Atoms of a given element are identical in size, mass and other properties and are different from atoms of other elements 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

Issues with Dalton’s Atomic Theory Atoms can be split into even smaller particles (nuclear chemistry) and aren’t indivisible i.e. nucleus, protons, electrons A given element can have different masses i.e. isotopes

J. J. Thomson (1897 A.D.) Discovered: The 1 st subatomic particle: the negatively charged electron, developed the plum pudding model Used a Cathode Ray Experiment Cathode Ray Tube – Electric current passed through a metal disk to another metal disk in a gas at low pressure (vacuum sealed tube) i. e. neon signs and ‘old-fashioned’ television sets

Cathode Ray Experiment When a current passed through the cathode ray tube, the surface of the tube opposite the cathode glowed Glow was hypothesized to be stream of particles called a cathode ray Ray affected by magnetic fields Attracted to positive charge Deflected from negative charge wMWcxeX8&NR=1 wMWcxeX8&NR=1

Discovery of the 1 st subatomic particle Thomson measured the ratio of the charge of the particles to their mass Same ratio no matter what metal or gas was used Named this particle an electron

J. J. Thomson’s Plum Pudding Model Atoms are electrically neutral Must have positive charges to balance the negatively charged electrons Electrons have a lot less mass than atoms Other particles must account for their mass Plum Pudding Model positively charged sphere with electrons dispersed through it

Ernest Rutherford (1911 A.D.) Discovered: The nucleus using the gold foil experiment Gold Foil Experiment Bombarded thin piece of gold foil with alpha particles Expected alpha particles to pass through with minimal deflection Surprised when 1 in 8000 deflected back to source It was “as if you had fired a 15 inch artillery shell at a piece of tissue paper and it came back and hit you”

Rutherford’s New Model of the Atom Discovered the nucleus is a small densely packed volume of positive charge Surrounded by cloud of electrons (more on this in chapter 4!) Size comparison Nucleus = marble Whole Atom = football field

Lesson 2.2: Basic Laws Learning Target: I will understand 3 of the basic laws of chemistry. Success Criteria: I can define the law of conservation of mass, the law of multiple proportions, and the law of definite proportions.

1790s – Discovery of Basic Laws Law of Conservation of Mass Mass is neither created nor destroyed during ordinary chemical reactions or physical changes Law of Definite Proportions A chemical compound contains the same elements in exactly the same proportions by mass regardless of size of sample or source of compound i.e. Every sample of table salt is made of 39.34% Na and 60.66% Cl i.e. H 2 O always has 2 atoms of H and 1 atom of O

Basic Laws Continued Law of Multiple Proportions If two or more different compounds are composed of the same two elements then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers i.e. CO and CO 2 CO = 1.00g of C and 1.33 g of O CO 2 = 1.00 g of C and 2.66 g of O The ratio of the second element is 2.66 to 1.33 or 2 to 1

Lesson 2.3: Inside the Atom Learning Target: I will understand the structure of the atom. Success Criteria: I can describe the location, mass, and charge of each subatomic particle.

Recap: Structure of the Atom Atom (today’s definition) = Smallest particle of an element that retains the chemical properties of that element Two regions: Nucleus Very dense, small center of the atoms Protons and neutrons Electron Cloud Region occupied by electrons

Inside the Nucleus 2 types of particles inside the nucleus: Protons = positively charged (+1) Neutrons = neutral charge (0) Mass in the nucleus Protons = 1 amu Neutrons = 1 amu 1 amu (atomic mass unit) = 1.66 x g

Where are the Electrons? In the Electron Cloud A cloud of negative charge outside of the nucleus (more in chapter 4!) Electrons = Negatively charged particles (-1) Mass of electrons = 0 amu Actually it is x g, but we round to zero for simplicity

Properties of Subatomic Particles ParticleSymbolChargeMass (amu) Electrone -, 0 e0 Protonp +, 1 H+11 Neutronn◦, 1 n01

Characteristics of Atoms Atomic Number = number of protons Identifies the element # of protons is what give that element its characteristic properties Elements with different protons are NOT THE SAME ELEMENT!!!

Lesson 3.4: Atoms, Ions, and Isotopes Learning Target: I will understand what atoms, ions, and isotopes are, and their relationship to subatomic particles. Success Criteria: I can count the number of subatomic particles of each type in atoms, isotopes, and ions.

Atoms Atoms have a neutral charge (charge=0) total positive charge equals the total negative charge # protons (+1 each) = # electrons (-1 each)

Ions Atoms with a charge Negative – more electrons than protons Positive – more protons than electrons Charge = #protons - # electrons Example: Magnesium atom with 12 protons and 10 electrons has a charge of +2

Isotopes Atoms of the same element (i.e. same # of protons) that have differing number of neutrons Isotopes of the same element: have different masses do not differ significantly in chemical behavior because they have the same # of protons Are named by their mass number

Mass Number Mass number = #protons + # neutrons ElementAtomic Number # of Protons # of Neutrons Mass Number Carbon66612 Oxygen88816 Nitrogen77815

Lesson 2.5: Average Atomic Mass Learning Target: I will understand the relationship of isotopes to an atoms average atomic mass. Success Criteria: I can calculate the average atomic mass of an element given the mass and percent abundance of each of it’s isotopes.

Average Atomic Mass Every element has isotopes The periodic table takes into account all naturally occurring isotopes of an element and averages them Average Atomic Mass is listed on the periodic table UNIT is amu = atomic mass unit How to Calculate Average Atomic Mass: Mass of isotope #1 x abundance in nature (decimal) + Mass of isotope #2 x abundance in nature (decimal) + … = Average Atomic Mass

Example of Calculating the Average Atomic Mass – Hydrogen There are two naturally occurring isotopes of hydrogen Hydrogen with 1 proton and zero neutrons Hydrogen with 1 proton and one neutron Differentiating between the two isotopes (symbol – mass number) Calculation: ( x ) + ( x ) = 1.01 amu Hydrogen IsotopesElement – mass #Atomic MassNaturally occurring abundance % 1 proton 0 neutrons H amu proton 1 neutron H amu0.0115