Prior Knowledge Check n Who first came up with the idea of the atom? n What particles are found in atoms? n Where are these particles found?

Section 5.1 Atoms n OBJECTIVES: –Summarize Dalton’s atomic theory. –Describe the size of an atom.

History of the atom n Not the history of atom, but of the idea of the atom. n Original idea Ancient Greece (400 B.C.) n There were 2 different ideas n Democritus and Leucippus- Greek philosophers.

History of Atom n Looked at beach n Made of sand n Cut sand - smaller sand n Smallest possible piece? n Atomos - not to be cut

Another Greek n Aristotle - Famous philosopher n All substances are made of 4 elements n Fire - Hot n Air - light n Earth - cool, heavy n Water - wet n Blend these in different proportions to get all substances

Who Was Right? n Greeks did not experiment. n Greeks settled disagreements by argument. n Aristotle was more famous. n He won. n His ideas carried through middle ages. n Alchemists change lead to gold.

Who’s Next? n Late 1700’s - John Dalton- England. n Teacher- summarized results of his experiments and those of others. n Dalton’s Atomic Theory n Combined ideas of elements with that of atoms.

Dalton’s Atomic Theory  All matter is made of tiny indivisible particles called atoms.  Atoms of the same element are identical, those of different elements are different.  Atoms of different elements combine in whole number ratios to form compounds.  Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

Just How Small Is an Atom? n An atom is the smallest particle of an element that retains the properties of that element n Atoms-very small: Fig. 5.2, p. 108 –still observable with proper instruments: Fig. 5.3, page 108

Section 5.2 Structure of the Nuclear Atom n OBJECTIVES: –Distinguish among protons, electrons, and neutrons in terms of relative mass and charge. –Describe the structure of an atom, including the location of the protons, electrons, and neutrons with respect to the nucleus. –Describe the experiments of Thomson and Rutherford.

Parts of Atoms n J. J. Thomson - English physicist n Made a piece of equipment called a cathode ray tube. n It is a vacuum tube - all the air has been pumped out.

Thomson’s Experiment Voltage source +- Vacuum tube Metal Disks

Thomson’s Experiment Voltage source +-

Thomson’s Experiment Voltage source +-

Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

Thomson’s Experiment n By adding an electric field

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field he found that the moving pieces were negative + -

Other particles n Proton - positively charged pieces 1840 times heavier than the electron – by E. Goldstein (~1886) n Neutron - no charge but the same mass as a proton – by J. Chadwick (1932) n Where are the pieces?

Rutherford’s experiment n Ernest Rutherford -English physicist. (1910) n Believed in the plum pudding model of the atom (discussed in Chapter 13). n Wanted to see how big they are. n Used radioactivity. n Alpha particles - positively charged pieces- helium atoms minus electrons n Shot them at gold foil which can be made a few atoms thick.

Plum Pudding- Yuck!

Rutherford’s experiment n When an alpha particle hits a fluorescent screen, it glows. n Here’s what it looked like (page 111)

Lead block Uranium Gold Foil Fluorescent Screen

He Expected n The alpha particles to pass through without changing direction very much. n Because…? n …the positive charges were thought to be spread out evenly. Alone they were not enough to stop the alpha particles.

What he expected

Because

He thought the mass was evenly distributed in the atom

Since he thought the mass was evenly distributed in the atom

What he got

How he explained it + n Atom is mostly empty. n Small dense, positive piece at center. n Alpha particles are deflected by it if they get close enough.

+

Density and the Atom n Since most of the particles went through, it was mostly empty space. n Because the pieces turned so much, the positive pieces were heavy. n Small volume, big mass, big density. n This small dense positive area is the nucleus.

Subatomic particles – p.111 Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ n0n / x x

Section 5.3 Distinguishing Between Atoms n OBJECTIVES: –Explain how the atomic number identifies an element.

Section 5.3 Distinguishing Between Atoms n OBJECTIVES: –Use the atomic number and mass number of an element to find the numbers of protons, electrons, and neutrons.

Section 5.3 Distinguishing Between Atoms n OBJECTIVES: –Explain how isotopes differ, and why the atomic masses of elements are not whole numbers.

Section 5.3 Distinguishing Between Atoms n OBJECTIVES: –Calculate the average atomic mass of an element from isotope data.

Counting the Pieces n Atomic Number = number of protons in the nucleus n # of protons determines kind of atom (since all protons are alike!) n the same as the number of electrons in the neutral atom. n Mass Number = the number of protons + neutrons. n These account for most of mass

Symbols n Contain the symbol of the element, the mass number and the atomic number.

Symbols X Mass number Atomic number

Symbols n Find the –number of protons –number of neutrons –number of electrons –Atomic number –Mass Number F 19 9

Symbols n Find the –number of protons –number of neutrons –number of electrons –Atomic number –Mass Number Br 80 35

Symbols n if an element has an atomic number of 34 and a mass number of 78 what is the –number of protons –number of neutrons –number of electrons –Complete symbol

Symbols n if an element has 91 protons and 140 neutrons what is the –Atomic number –Mass number –number of electrons –Complete symbol

Symbols n if an element has 78 electrons and 117 neutrons what is the –Atomic number –Mass number –number of protons –Complete symbol

Isotopes n Dalton was wrong. n Atoms of the same element can have different numbers of neutrons. n different mass numbers. n called isotopes.

Naming Isotopes n We can also put the mass number after the name of the element. n carbon- 12 n carbon -14 n uranium-235

Atomic Mass n How heavy is an atom of oxygen? –There are different kinds of oxygen atoms. n More concerned with average atomic mass. n Based on abundance of each element in nature. n Don’t use grams because the numbers would be too small.

Measuring Atomic Mass n Unit is the Atomic Mass Unit (amu) n One twelfth the mass of a carbon- 12 atom. n Each isotope has its own atomic mass, thus we determine the average from percent abundance.

Calculating averages n Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. n Sample 5-5, p.120

Atomic Mass n Calculate the atomic mass of copper if copper has two isotopes % has a mass of 63 amu and the rest has a mass of 64 amu.

Atomic Mass n Magnesium has three isotopes % magnesium 24, 10.00% magnesium 25, and the rest magnesium 26. What is the atomic mass of magnesium? n If not told otherwise, the mass of the isotope is the mass number in amu

Atomic Mass n Is not a whole number because it is an average. n are the decimal numbers on the periodic table.

Section 5.4 The Periodic Table: Organizing the Elements n OBJECTIVES: –Describe the origin of the periodic table.

Section 5.4 The Periodic Table: Organizing the Elements n OBJECTIVES: –Identify the position of groups, periods, and the transition metals in the periodic table.

Development of the Periodic Table n mid-1800s, about 70 elements n Dmitri Mendeleev – Russian chemist n Arranged elements in order of increasing atomic mass n Thus, the first “Periodic Table”

Mendeleev n Left blanks for undiscovered elements n When discovered, good prediction n Problems? –Co and Ni; Ar and K; Te and I

New way n Henry Moseley – British physicist n Arranged elements according to increasing atomic number n The arrangement today n P.124 – long form n Symbol, atomic number & mass

Periodic table n Horizontal rows = periods –There are 7 periods n Periodic law: when the elements are arranged in order of atomic number, then there are periodic repetition of physical and chemical properties n Vertical column = group (or family) –Similar physical & chemical prop. –Identified by number & letter

Areas of the periodic table n Group A elements = representative elements –Wide range of phys & chem prop. –Metals, nonmetals, metalloids (semimetals) –Very reactive, not reactive –Solids, liquids and gases

Metals n Group IA – alkali metals n Group 2A – alkaline earth metals n Transition metals and Inner transition metals – Group B n Metals: electrical conductors, have luster, ductile, malleable n All metals are solids at room temperature, except _____.

Nonmetals n Nonmetals: generally nonlustrous, poor conductors of electricity –Some gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) n Group 7A – halogens n Group 0 – noble gases

Division between metal & nonmetal n Heavy, stair-step line n Metalloids border the line –Properties intermediate between metals and nonmetals n Learn the general behavior and trends of the elements, instead of memorizing each element property