chapter 2 atoms and ions -...

Chapter 2

Atoms and Ions

A History of Atomic Models

400 B.C.E.(Democritus,

a early atomist)

1804 (Dalton)

Law of Conservation of Mass

Antoine Lavoisier1743-1794

In a chemical reaction, matter is neither created nor destroyed. Total mass of the materials you have before the reaction must

equal the total mass of the materials you have at the end.

total mass of reactants = total mass of products

7.7 g Na + 11.9 g Cl2 19.6 g NaCl

Law of Conservation of Mass

Law of Definite Proportions

Joseph Proust1754-1826

All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements.

Proportions of Na and Cl in Sodium Chloride

A 100.0 g sample of sodium chloride contains 39.3 g of sodium and 60.7 g of chlorine



Law of Multiple ProportionsJohn Dalton1766-1844

When two elements (call them A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small, whole numbers.

Carbon monoxide contains 1.33 g of oxygen for every 1.00 g of carbon.

Carbon dioxide contains 2.67 g of oxygen for every 1.00 g of carbon.

Law of Multiple Proportions

1.33/1.00 = 4/3 2.67/1.00 = 8/3

Dalton’s Atomic Theory

1. Each element is composed of tiny, indestructible particles called atoms.

2. All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements.

3. Atoms combine in simple, whole-number ratios to form molecules of compounds.

4. In a chemical reaction, atoms are reorganized in the way they are bound together.

5. Atoms of one element cannot change into atoms of another element.

Some Notes on Charge

Two kinds of charge called + and –

Opposite charges attract+ attracted to –

Like charges repel+ repels +– repels –

To be neutral, something must have no charge or equal amounts of opposite charges

J.J. Thomson and Cathode Rays

Examined the electrically charge “particles” in cathode raysMeasured the amount of force necessary to deflect the paths of

the particles

The particles have a negative charge. The amount of deflection was related to two factors, the charge

and mass of the particles.

Every material tested contained these same particles.

+++++++++++

-------------

Power Supply - +

Cathode

(-) (+)

Thomson’s Conclusions

If the particle has the same amount of charge as a hydrogen ion, then it must have a mass almost 2000 x smaller than hydrogen atoms!

The only way for this to be true is if these particles were pieces of atoms.

Apparently, the atom is not unbreakable.

Thomson believed that these particles were the ultimate building blocks of matter.

These cathode ray particles became known as electrons.

A New Theory of the Atom

Thomson proposed that instead of being a hard, marble-like unbreakable sphere, the way Dalton described it, the atom actually had an inner structure.

The structure of the atom contains many negatively charged electrons. These electrons are held in the atom by their attraction for a positively charged electric field within the atom.



a early atomist)

1804 (Dalton) 1903

(Thomson)

Predictions of the Plum Pudding Atom

The mass of the atom is due to the mass of the electrons within it. Electrons are the only particles in Plum Pudding atoms, therefore the only source of mass.

The atom is mostly empty space and should not have a bunch of negatively charged particles near each other as they would repel.

Radioactivity

In the late 1800s, Henri Becquerel and Marie Curie discovered that certain elements would constantly emit

small, energetic particles and rays.

These energetic particles could penetrate matter.

Ernest Rutherford discovered that there were three different kinds of emissions:

alpha rays made of particles with a mass 4 x H atom and + charge

beta rays made of particles with a mass ~1/2000th H atom and – charge

gamma rays that are energy rays, not particles

Explaining Results of α-particle Scattering Experiments

Rutherford’s expectation

Small, positively charged α-particles should pass through the nebulous, positively charged cloud of the Thomson atomic model largely undeflected. Some would be slightly deflected by passing near electrons (present to neutralize the positive charge of the cloud).

Actual Result

To minimize alpha loss by scattering from air molecules, the experiment was carried out in a fairly good vacuum, the metal box being evacuated through a tube T (see below). The alphas came from a few milligrams of radium (to be precise, its decay product radon 222) at R in the figure below, from the original paper, which goes on: "By means of a diaphragm placed at D, a pencil of alpha particles was directed normally on to the scattering foil F. By rotating the microscope [M] the alpha particles scattered in different directions could be observed on the screen S." Actually, this was more difficult than it sounds. A single alpha caused a slight fluorescence on the zinc sulphide screen S at the end of the microscope. This could only be reliably seen by dark-adapted eyes (after half an hour in complete darkness) and one person could only count the flashes accurately for one minute before needing a break, and counts above 90 per minute were too fast for reliability. The experiment accumulated data from hundreds of thousands of flashes.

Rutherford's partner in the initial phase of this work was Hans Geiger, who later developed the Geiger counter to detect and count fast particles. Many hours of staring at the tiny zinc sulphide screen in the dark must have focused his mind on finding a better way!

In 1909, an undergraduate, Ernest Marsden, was being trained by Geiger. To quote Rutherford (a lecture he gave much later):

"I had observed the scattering of alpha-particles, and Dr. Geiger in my laboratory had examined it in detail. He found, in thin pieces of heavy metal, that the scattering was usually small, of the order of one degree. One day Geiger came to me and said, "Don't you think that young Marsden, whom I am training in radioactive methods, ought to begin a small research?" Now I had thought that, too, so I said, " Why not let him see if any alpha-particles can be scattered through a large angle?" I may tell you in confidence that I did not believe that they would be, since we knew the alpha-particle was a very fast, massive particle with a great deal of energy, and you could show that if the scattering was due to the accumulated effect of a number of small scatterings, the chance of an alpha-particle's being scattered backward was very small. Then I remember two or three days later Geiger coming to me in great excitement and saying "We have

been able to get some of the alpha-particles coming backward …" It was quite the most incredible event that ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you."

Rutherford’s Results

Over 98% of the α particles went straight through

About 2% of the α particles went through but were deflected by large angles

About 0.005% of the α particles bounced off the gold foil

“...as if you fired a 15” cannon shell at a piece of tissue paper and it came back and hit you.”

The atom contains a tiny dense center called the nucleus.The amount of space taken by the nucleus is only about 1/109

the volume of the atom.

The nucleus has essentially the entire mass of the atom.The electrons weigh so little they give practically no mass to the

atom.

The nucleus is positively charged. The amount of positive charge balances the negative charge of

the electrons.

The electrons are dispersed in the empty space of the atom surrounding the nucleus.

Rutherford’s Interpretation – the Nuclear Model

Structure of the Nucleus

Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron but opposite sign – the proton.

Because protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons.

Some Problems

How could beryllium have four protons stuck together in the nucleus?

If a beryllium atom has four protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from?

Protons - subatomic particles found in the nucleus

The proton has a charge of +1.60 x 1019 C. The proton has a mass of 1.67262 x 10−24 g.

Electrons - negatively charged particles found in all atoms. Cathode rays are made of streams of electrons. The electron has a charge of −1.60 x 1019 C. The electron has a mass of 9.1 x 10−28 g.

There Must Be Something Else!

To answer these questions, Rutherford and Chadwick proposed that there was another particle in the nucleus – it is called a neutron.

Neutrons are subatomic particles with a mass = 1.67493 x 10−24 g and no charge, and are found in the nucleus



a early atomist)

1804 (Dalton) 1903

(Thomson)

1911 (Rutherford)1913

(Bohr)1932

(Chadwick)

1 amu = 1 atomic mass unit

Quantum Model

The Number of Protons Defines the Element

Structure of the Nucleus

Soddy (1913) discovered that the same element could have atoms with different masses, which he called isotopes.

The observed atomic mass is a weighted average of the weights of all the naturally occurring atoms.

The percentage of an element that is one isotope is called the isotope’s natural abundance.

Atomic Mass (Weight)

Carbon 98.99 % C-12

1.11 % C-13 Weighted average = 12.01 amu trace % C-14 }

Chlorine 75.5 % Cl-35 Weighted average = 35.45 amu 24.5 % Cl 37 }

Isotopes

All isotopes of an element are chemically identical and undergo the exact same chemical reactions.

All isotopes of an element have the same number of protons.

Isotopes of an element have different masses.

Isotopes of an element have different numbers of neutrons.

Isotopes are identified by their mass numbers, which is the sum of all the protons and neutrons in the nucleus.

Isotopes

Atomic numberNumber of protons = Z

Mass Number

Protons + neutrons = whole number = A

Relative Abundance = relative amount found in a sample

XZ

A

# neutrons = A-Z

Al27

13

C13

6

Mo96

42

Cs133

55

Complete the table:

6 6 13

42 54 42

13 14 13 13 27

55 78 55

Charged Atoms

Ions are atoms which have acquired a charge through loss or gain of electrons.

When atoms gain electrons, they become negatively charged ions, called anions.

When atoms lose electrons, they become positively charged ions, called cations.

Complete the table

+3Al

16 2-

12 10

13 13 10 3+

35 35

S2-

Mg2+

Br-

Periodic Trend in Ion Formation

Nonmetals form anions.

Anions named by changing the ending of the name to -ide.

F + 1e− F−

fluorine fluoride atom ion

O + 2e− O2−

oxygen oxide atom ion

Metals form cations.

Cations are named the same as the metal.

Na Na+ + 1e− sodium sodium atom ion

Ca Ca2+ + 2e− calcium calcium atom ion

Periodic Trend in Ion Formation

Main-group ions and the noble gas configurations.

Nonmetals Metals

chapter 2 atoms and ions -...

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