87
Atoms: The Building Blocks of Matter 1
Atoms: The Building Blocks
of Matter
1
WHEN YOU CRUSH A LUMP OF SUGAR, YOU CAN SEE THAT IT IS MADE UP OF MANY SMALLER PIECES OF SUGAR. YOU MAY GRIND THESE PARTICLES INTO A VERY FINE POWDER, BUT EACH TINY PIECE IS STILL SUGAR.
The Atom: From Philosophical Idea to Scientific Theory
2
3
NOW SUPPOSE YOU DISSOLVE THE SUGAR IN WATER. THE TINY PARTICLES SEEM TO DISAPPEAR COMPLETELY.
4
EVEN IF YOU LOOK AT THE SUGAR-WATER SOLUTION THROUGH A POWERFUL MICROSCOPE YOU CANNOT SEE ANY SUGAR PARTICLES. YET IF YOU WERE TO TASTE THE SOLUTION, YOU’D KNOW THAT THE SUGAR IS STILL THERE.
5
OBSERVATIONS LIKE THESE LED EARLY PHILOSOPHERS TO PONDER THE FUNDAMENTAL NATURE OF MATTER. IS IT CONTINUOUS AND INFINITELY DIVISIBLE, OR IS IT DIVISIBLE ONLY UNTIL A BASIC, INVISIBLE PARTICLE THAT CANNOT BE DIVIDED FURTHER IS REACHED?
Foundations of Atomic Theory
Nearly all chemists in late 1700s accepted the definition of an element as a substance that cannot be broken down further
6
7
They knew about chemical reactions but there was great disagreement as to whether elements always combine in the same ratio when forming a specific compound
Law of Conservation of
Mass
With the help of improved balances, investigators could accurately measure the masses of the elements and compounds they were studying
8
9
This lead to discovery of several basic laws
Law of conservation of mass - states that mass is neither destroyed nor created during ordinary chemical reactions or physical changes
Law of Definite Proportions
law of definite proportions - A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound
10
11
EACH OF THE SALT CRYSTALS SHOWN HERE CONTAINS
EXACTLY 39.34% SODIUM AND 60.66% CHLORINE BY
MASS.
Law of Multiple Proportions
Two elements sometimes combine to form more than one compound
For example, the elements carbon and oxygen form two compounds, carbon dioxide and carbon monoxide
Consider samples of each of these compounds, each containing 1.0 g of carbon
12
13
In carbon dioxide, 2.66 g of oxygen combine with 1.0 g of carbon
In carbon monoxide, 1.33 g of oxygen combine with 1.0 g of carbon The ratio of the masses of oxygen in these two compounds is exactly 2.66 to 1.33, or 2 to 1
1808 John Dalton
Proposed an explanation for the law of conservation of mass, the law of definite proportions, and the law of multiple proportions He reasoned that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds
14
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; atoms of different elements differ in size, mass, and other properties.
15
3. Atoms cannot be divided, 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.
16
Modern Atomic Theory
Dalton turned Democritus’s idea into a scientific theory which was testable Not all parts of his theory have been proven correct
17
18
Ex. We know atoms are divisible into even smaller particles - subatomic particles
19
We know an element can have atoms with different masses - isotopes
Isotone
different atoms with the same number of neutrons
20
how to remember?
isotoPe - same number of Protons isotoNe - same number of Neutrons
21
isobar
different elements, same atomic mass
22
ALTHOUGH JOHN DALTON THOUGHT ATOMS WERE INDIVISIBLE, INVESTIGATORS IN THE LATE 1800S PROVED OTHERWISE. AS SCIENTIFIC ADVANCES ALLOWED A DEEPER EXPLORATION OF MATTER, IT BECAME CLEAR THAT ATOMS ARE ACTUALLY COMPOSED OF SEVERAL BASIC TYPES OF SMALLER PARTICLES AND THAT THE NUMBER AND ARRANGEMENT OF THESE PARTICLES WITHIN AN ATOM DETERMINE THAT ATOM’S CHEMICAL PROPERTIES.
The Structure of the Atom23
24
TODAY WE DEFINE AN ATOM AS THE SMALLEST PARTICLE OF AN ELEMENT THAT RETAINS THE CHEMICAL PROPERTIES OF THAT ELEMENT.
All atoms consist of two regions
Nucleus - very small region located near the center of an atom In the nucleus there is at least one positively charged particle called the proton
25
26
Usually at least one neutral particle called the neutron Surrounding the nucleus is a region occupied by negatively charged particles called electrons
Discovery of the Electron
Resulted from investigations into the relationship between electricity and matter
Late 1800s, many experiments were performed: electric current was passed through different gases at low pressures
27
28
These experiments were carried out in glass tubes known as cathode-ray tubes
Cathode Rays and Electrons
Investigators noticed that when current was passed through a cathode-ray tube, the opposite end of the tube glowed
29
30
Hypothesized that the glow was caused by a stream of particles, which they called a cathode ray The ray traveled from the cathode to the anode when current was passed through the tube
Observations
1. cathode rays deflected by magnetic field in same way as wire carrying electric current (known to have negative charge) 2. rays deflected away from negatively charged object
31
32
Observations led to the hypothesis that the particles that compose cathode rays are negatively charged Strongly supported by a series of experiments carried out in 1897 by the English physicist Joseph John Thomson
33
He was able to measure the ratio of the charge of cathode-ray particles to their mass He found that this ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube
34
Thomson concluded that all cathode rays are composed of identical negatively charged particles, which were later named electrons
35
Charge and Mass of the Electron
Confirmed that the electron carries a negative electric charge Because cathode rays have identical properties regardless of the element used to produce them, it was concluded that electrons are present in atoms of all elements
36
Cathode-ray experiments provided evidence that atoms are divisible and that one of the atom’s basic constituents is the negatively charged electron
37
Thomson’s experiment revealed that the electron has a very large charge for its tiny mass
Mass of the electron is about one two-thousandth the mass of the simplest type of hydrogen atom (the smallest atom known) Since then found that the electron has a mass of 9.109 × 10−31 kg, or 1/1837 the mass of the hydrogen atom 38
Based on information about electrons, two inferences made about atomic structure
b/c atoms are neutral they must have positive charge to balance negative electrons
b/c electrons have very little mass, atoms must have some other particles that make up most of the mass
39
Thomson’s Atom
Plum pudding model (based on English dessert)
Negative electrons spread evenly through positive charge of the rest of the atom Like seeds in a watermelon
40
Discovery of Atomic Nucleus
1911 by New Zealander Ernest Rutherford and his associates Hans Geiger and Ernest Marsden
Bombarded a thin, gold foil with fast-moving alpha particles (positively charged particles with about four times the mass of a hydrogen atom)
41
Assume mass and charge were uniformly distributed throughout atoms of gold foil (from Thomson’s model of the atom) Expected alpha particles to pass through with only slight deflection
42
What Really Happened…
Most particles passed with only slight deflection
However, 1/8,000 were found to have a wide deflection
43
Rutherford explained later it was “as if you have fired a 15-inch artillery shell at a piece of tissue paper and it came back and hit you.”
44
45
Explanation
After 2 years, Rutherford finally came up with an explanation The rebounded alpha particles must have experienced some powerful force within the atom
46
47
The source of this force must occupy a very small amount of space because so few of the total number of alpha particles had been affected by it The force must be caused by a very densely packed bundle of matter with a positive electric charge Rutherford called this positive bundle of matter the nucleus
Rutherford had discovered that the volume of a nucleus was very small compared with the total volume of an atom
48
If the nucleus were the size of a marble, then the size of the atom would be about the size of a football field 49
50
But where were the electrons? Rutherford suggested that the electrons surrounded the positively charged nucleus like planets around the sun
He could not explain, however, what kept the electrons in motion around the nucleus
Rutherford’s Atom
51
Composition of Atomic Nucleus
Except hydrogen, all atomic nuclei made of two kinds of particles Protons Neutrons
52
53
Protons = positive Neutrons = neutral Electrons = negative
Atoms are electrically neutral, so number of protons and electrons IS ALWAYS THE SAME
The nuclei of atoms of different elements differ in the number of protons they contain and therefore in the amount of positive charge they possess
So the number of protons in an atom’s nucleus determines that atom’s identity
54
Forces in the Nucleus
Usually, particles that have the same electric charge repel one another Would expect a nucleus with more than one proton to be unstable
55
When two protons are extremely close to each other, there is a strong attraction between them Nuclear forces - short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together
56
The Sizes of atoms
Area occupied by electrons is electron cloud – cloud of negative charge Radius of atom is distance from center of nucleus to outer portion of cloud
57
Unit – picometer (10-12 m) Atomic radii range from 40-270 pm Very high densities – 2 x 108 tons/cm3
58
Counting Atoms
Consider Neon, Ne, the gas used in many illuminated signs. Neon is a minor part of the atmosphere. In fact, dry air contains only about 0.002% Ne. And yet there are about 5 x 1017 atoms of neon present in each breath you inhale.
59
In most experiments, atoms are too small to be measured individually. Chemists can analyze atoms quantitatively, however, by knowing fundamental properties of the atoms of each element.
60
In this section, you will be introduced to some of the basic properties of atoms. You will then discover how to use this information to count the number of atoms of an element in a sample with a known mass.
61
You will also become familiar the the mole, a special unit used by chemists to express amounts of particles, such as atoms and molecules.
62
Atomic Number
Atoms of different elements have different numbers of protons
Atomic number (Z) - number of protons in the nucleus of each atom of that element
63
64
Shown on periodic table Atomic number identifies an element
65
Mass Number
Mass number - total number of protons and neutrons in the nucleus of an isotope 66
Shows the composition of a nucleus as the isotope’s nuclear symbol Uranium-235 is written as
Nuclide – general term for specific isotope of an element
67
23592U
68
The superscript indicates the mass number
69
the subscript indicates the atomic number
70
Relating Mass to Numbers of Atoms
The relative atomic mass scale makes it possible to know how many atoms of an element are present in a sample of the element with a measurable mass
71
72
Three very important concepts provide the basis for relating masses in grams to numbers of atoms 1. The mole
2. Avogadro’s number 3. Molar mass
The Mole
SI unit for an amount of substance
(like 1 dozen = 12) Mole (mol) - amount of a substance that contains as many particles as there are atoms in exactly 12g of C-12
73
Avogadro’s Number
Avogadro’s number - the number of particles in exactly one mole of a pure substance
6.022 x 1023 How big is that?
74
75
If 5 billion people worked to count the atoms in one mole of an element, and if each person counted continuously at a rate of one atom per second, it would take about 4 million years for all the atoms to be counted
Molar Mass
Molar mass - the mass of one mole of a pure substance
Written in unit g/mol
Found on periodic table (atomic mass)
Ex. Molar mass of H = 1.008 g/mol76
2.00 mol He x 4.00 g He = 8.00 g He1 mole He
Example
How many grams of helium are there in 2 moles of helium?
2.00 mol He x = ? g He
77
Practice Problems
What is the mass in grams of 3.50 mol of the element copper, Cu? 222g Cu
78
79
What is the mass in grams of 2.25 mol of the element iron, Fe? 126g Fe What is the mass in grams of 0.375 mol of the element potassium, K? 14.7g K
80
What is the mass in grams of 0.0135 mol of the element sodium, Na? 0.310g Na What is the mass in grams of 16.3 mol of the element nickel, Ni? 957g Ni
1. A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced? 0.441 mol Al
2. How many moles of calcium, Ca, are in 5.00 g of calcium? 0.125 mol Ca
3. How many moles of gold,Au, are in 3.60 × 10−10g of gold?
1.83 × 10−12 mol Au81
Conversions with Avogadro’s Number
How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?
Given: 3.01x1023 atoms of Ag Unknown: amount of Ag in moles
Ag atoms × = moles Ag
82
83
3.01x1023atomsAg x 1 mol Ag =0.500 mol Ag6.02x1023atomsAg
Practice problems
1. How many moles of lead, Pb, are in 1.50×1012 atoms of lead?
2.49×10−12 mol Pb
84
85
2. How many moles of tin, Sn, are in 2500 atoms of tin?
4.2×10−21 mol Sn
86
3. How many atoms of aluminum, Al, are in 2.75 mol of aluminum?
1.66×1024 atoms Al
1. What is the mass in grams of 7.5 × 1015 atoms of nickel, Ni? 7.3×10−7g Ni 2. How many atoms of sulfur, S, are in 4.00 g of sulfur?
7.51×1022atoms S 3. What mass of gold,Au, contains the same number of atoms as 9.0 g of aluminum,Al? 66g Au
87
