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Atoms and ElementsChapter 2

Some images Copyright © The McGraw-Hill Companies, Inc.

ed. Brad Collins

Sunday, August 18, 13

Early Ideas of the atom

•5th Century BC - Democritus

• All matter consists of very small, indivisible particles.

• Called them atmos


Alchemists• Interested in transmutation of elements

• Mystical aspects - Philosopher's Stone

• Developed many techniques still in use:

• Distillation

• Purification by crystallization

• Some of the first users of the Scientific Method

• Abu Musa Jābir ibn Hayyān (Geber)

• Paracelsus (Philippus Theophrastus Aureolus Bombastus von Hohenheim)

• "The dose makes the poison"

• 'Father' of toxicology

Geber, ca. 721-815

Paracelsus, 1492-1541


Law of Conservation of Mass• Roots in Ancient Greek and Indian (Jain) philosophy

• Empedocles ca.450 BCE "Nothing comes from nothing"

• Mahavira ca.500 BCE 'The universe and its constituents cannot be destroyed or created'.

• Nasīr al-Dīn al-Tūsī ca.1250 AD

• "A body of matter cannot disappear completely. It only changes its form, condition, composition, color and other properties and turns into a different complex or elementary matter".

• Mikhail Lomonosov (1748) showed conservation so mass experimentally

• Antoine Lavoisier (1774) "Matter cannot be created or destroyed in a chemical reaction".

• Mercury (II) oxide (solid) + heat ---> Mercury (liquid) + Oxygen (gas)

• 50.0 g 46.3 g 3.7 g


Law of Definite Proportions• 1799 - Joseph Proust

• Different samples of the same compound always contain its constituent elements in the same proportion by mass.

• Lead sulfide combines in a ratio of 6.5 g lead to 1 g sulfur.

• Doesn't matter how much of either you start with: 10 g of lead will react with 1 g of sulfur, leaving 3.5 g of unreacted lead behind.

• Works because atoms of a particular element have the same mass.

• Lead (Pb) 207.2, Sulfur (S) 32.06

• Lead sulfide contains one lead atom and one sulfur atom,

• 207.2:32.06 (6.46:1)


Law of Multiple Proportions• Elements can combine in more than one fixed mass ratio to form

different compounds

• Example: carbon and oxygen

• Can combine in a 3:4 ratio and 3:8 ratio

• 3 g carbon + 4 g oxygen make 7 g carbon monoxide

• 3 g carbon + 8 g oxygen make 11 g carbon dioxide

• Note that ratio of the oxygen masses is 2:1

• Law of multiple proportions says that ratio of elements in multiple compounds will be simple whole numbers


Dalton’s Atomic Theory (1808)1. Elements are composed of extremely small particles called

atoms. 2. All atoms of a given element are identical, having the same

size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements.

3. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same (fixed ratio).

4. Chemical reactions involve only the rearrangement of the atoms. Atoms are not created or destroyed in the chemical reactions.

The idea that atoms are neither created nor destroyed in chemical reactions is known as the Law of the Conservation of Mass. 2.1


Application of Dalton’s Laws • Dalton's Laws helped to explain several phenomena

associated with chemical compounds.• The Law of Definite Proportions:

• When atoms combine to form compounds they combine in a fixed ratio.

• Atoms of a given element all have the same weight.

Water, 2H + 1O Nitrogen, 2N Ammonia, 3H + 1N


Application of Dalton’s Laws • The Law of Multiple Proportions

• Definition: If two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

• Why?• Different compounds made up of the same elements differ in the

number of atoms of each kind that combine.Sunday, August 18, 13

Law of Multiple Proportions

2.1

Mass1612

Mass3212


8 X2Y16 X 8 Y+

Law of Conservation of Mass

2.1Sunday, August 18, 13

Atomic StructureDefinition: Atom - the basic unit of an element that can enter into chemical combination.

Dalton believed atoms were indivisible.Experiments beginning ca. 1850 showed atoms have internal structure.

Radiation: The emission and transmission of energy through space in the form of waves.

2.3


(Uranium compound)Radioactivity: the spontaneous emission of particles and/or radiation 2.3


J.J. Thomson, measured mass/charge of e- (1906 Nobel Prize in Physics)

2.3


Cathode Ray Tube

2.3


Miliken’s experiment: e- charge = -1.602 x 10-22 CJ.J. Thomson’s charge/mass of e- = -1.76 x 108 C/g e- mass = 9.10 x 10-28 g

Measured mass of e- (1923 Nobel Prize in Physics)

2.3


2.3


1. Atoms positive charge is concentrated in the nucleus2. Proton (p) has opposite (+) charge of electron (–) = +1.602 x 10–22 C3. Mass of p is 1.67 x 10-24 g, which is 1840 x mass of e–

4. Relative charge of proton is +1; electron is –1

α particle velocity ~ 1.4 x 107 m/s(~5% speed of light)

(1908 Nobel Prize in Chemistry)

2.3


Rutherford’s Atomic Model

Atom’s Positive Charge

Atom’s MassElectron Cloud

Question: If all the protons are contained in the nucleus, why don’t they repel each other?

Nucleus

2.3


An ion is an atom, or group of atoms, that has a net positive or negative charge.cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation.

anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion.

Na 11 protons11 electrons Na+ 11 protons

10 electrons

Cl 17 protons17 electrons

Cl-17 protons18 electrons

2.3


NeutronsObservations:

H atoms have 1 p; He atoms have 2 pMass He/Mass H should = 2

BUT: measured mass He/mass H = 4

Walter Bothe and Herbert Becker (1930):Bombarded beryllium with alpha particles gave off electrically neutral radiationGamma ray?

Frederic Joliot-Curie (1932):Experiments with Bothe-Becker radiation showed that if it fell on paraffin, it could eject protons with very high energy.

Could not be a gamma ray - photons have no massJames Chadwick (1932):

Showed that Bothe-Becker radiation could eject protons when it fell on substances other than paraffin.Energies of the protons ejected could be accounted for, if mass of neutral particle = mass of proton

Rutherford hypothesized that a neutral third particle must exist.

2.3


NeutronsNeutrons (n) are neutral (charge = 0)n mass ~ p mass = 1.67 x 10-24 g = 1 amu*e– mass = 5.45 x 10–4 amu

3.12.3

*amu = Atomic Mass Unit or Dalton (Da)


atomic radius ~ 100 pm = 1 x 10-10 mnuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m

Modern Model of the Atom

“If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”

2.3


Particle Mass (g) Mass (amu) Charge (C) Charge Unit

Electron 9.10939 x 10–28 5.45 x 10–4 –1.6022 x 10–19 –1

Proton 1.67262 x 10–24 1 +1.6022 x 10–19 +1

Neutron 1.67493 x 10–24 1 0 0

Summary: Subatomic Particle Properties

2.3


Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons

Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei

XAZ

H11 H (D)2

1 H (T)31

U23592 U238

92

Mass Number Atomic Number

Element Symbol

2.3


2.3


6 protons, 8 (14 - 6) neutrons, 6 electrons

6 protons, 5 (11 - 6) neutrons, 6 electrons

Do You Understand Isotopes?

How many protons, neutrons, and electrons are in C146 ?

How many protons, neutrons, and electrons are in C116 ?

2.3


Development of the Periodic Table• Dmitri Mendeleev (1834-1908)

• Saw a pattern of chemical properties in the known elements.

• Na similar to Li

• Mg similar to Be

• Made a set of cards that could be arranged based on this pattern (periodicity)

• Predicted that gaps in his table would be occupied by as yet undiscovered elements

• Gallium discovered in 1869, matched Mendeleev’s predicted properties

2.4


Period

Group

2.4


Other Forms of the Periodic Table

2.4


Atomic Mass• The weighted average of masses of the naturally occurring isotopes

of an element.

• Carbon-12 is defined as 12.00 amu

• All other elements masses are relative to carbon-12

• Chlorine has 2 naturally occurring isotopes:

• chlorine-35 (mass 34.97 amu) 75.77%

• chlorine-37 (mass 36.97 amu) 24.23%

• Atomic mass of chlorine = (34.97)(0.7577) + (36.97)(0.2433) = 35.45

2.4


Periodic TableA Group:

Main group elements

Representative elements

Inner Transition Elements:Sometimes called ‘Rare earth elements’

B Group:Transition elementsTransition metals

A Group:Main group elements

Representative elements

2.4


Periodic TableAlkali metals

Alkaline earth metals HalogensNoble Gases

2.4


Periodic TableMetals:

Good conductors of heat and electricity

Nonmetals:

Malleable, ductileHigh luster solids (except mercury)Lose electrons easily

Poor conductors of heat and electricityBrittleMay be gases, liquids or solids

Tend to gain electrons

Metaloids:Properties between metals and nonmetals e.g. fair conductor, but brittle.

2.4


Elements• 118 Known elements

• Elements are represented by a 1 or 2 character symbol.

• Sometimes based on English names:

• Hydrogen (H), Carbon (C)

• Helium (He), Cobalt (Co)

• Sometimes based on another language (e.g., Latin)

• Copper (Cu), Latin, Cuprum

• Tin (Sn) Latin, Stannum

• Mercury (Hg), Greek, Hydragyrum

• Tungsten (W), Sweedish, Wolfram

• Most elements are monatomic (exist as individual atoms)

• Some elements are diatomic (2-atom pairs): H2, N2, O2

• A few elements are polyatomic (groups of 3 or more atoms): sulfur (S8), phosphorus (P4)2.4


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