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Chapter 3A
Atoms and Elements
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CHAPTER OUTLINE���
§ Elements and Symbols § Periodic Table of the Elements § Properties of Metals and Non-Metals § The Atomic Theory § The Modern Atom § Atomic Structure § Isotopes and Atomic Mass
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ELEMENTS���AND SYMBOLS���
q Elements are primary substances from which all other substances are built. Elements cannot be broken down into simpler substances.
q Over time some elements have been named for planets, mythological figures, minerals, colors, geographic locations and famous people. Some examples are shown below:
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ELEMENTS���AND SYMBOLS���
q The symbol for most elements is the one- or two-letter abbreviation of the name of the element. Only the first letter of an elements symbol is capitalized. If the symbol has a second letter, it is written as lowercase.
Co cobalt
CO carbon and oxygen
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ELEMENTS���AND SYMBOLS���
q Although most of the symbols use letters from current names, some of the symbols of the elements are based on their Greek or Latin names.
Na sodium
Fe iron
(natrium)
(ferrum)
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ELEMENTS���AND SYMBOLS���
Hydrogen (H2)
Oxygen (O2)
Nitrogen (N2)
Fluorine (F2)
Chlorine (Cl2)
Bromine (Br2)
Iodine (I2)
q Some elements have formulas that are not single atoms. Seven of these elements have diatomic (2-atoms) molecules.
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PERIODIC TABLE���
Metals Non-metals Metalloids
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PROPERTIES OF ���METALS & NON-METALS���
Metals Non-metals
• Mostly solid • Can be solid, liquid or gas
• Have shiny appearance • Have dull appearance
• Good conductors of heat & electricity
• Poor conductors of heat & electricity
• Malleable & ductile • Brittle (if solid)
• Lose electrons • Gain or share electrons
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METALLOIDS���
q Metalloids are elements that possess some properties of metals and some of non-metals.
q The most important metalloids are silicon (Si) and germanium (Ge) which are used extensively in computer chips.
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PERIODIC TABLE���
Metallic character increases going down a group Metallic character decreases going across a period.
Cs Fr
Most metallic elements F
Least metallic element
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PERIODIC TABLE���
q Seven elements exist as diatomic molecules. q All others exist as monatomic (single atom).
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PERIODS & GROUPS���
q The periodic table is composed of periods (rows) and groups or families (columns).
q Elements in the same family have similar properties, and are commonly referred to by their traditional names.
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PERIODS & GROUPS���
q Elements in groups 1-2 and 13-18 are referred to as main-group or representative groups.
q Alkali metals are soft metals that are very reactive. They often react explosively with other elements.
q Noble gases are un-reactive gases that are commonly used in light bulbs.
q Halogens are the most reactive nonmetals, and occur in nature only as compounds.
q Group 2 elements are called alkaline-earth metals. These metals are less reactive than alkali metals.
q The group of metals in between the main group elements are called transition metals.
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EARLY CONCEPTS���OF THE ATOM���
q The smallest particle of matter that still retains its properties is called an atom.
q In the fifth century B.C., the Greek philosopher Democritus proposed that matter is composed of a finite number of discrete particles, named atomos (meaning un-cuttable or indivisible)
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DALTON’S ���ATOMIC THEORY���
q In 1808, John Dalton, built on ideas of Democritus, and formulated a precise definition of the building blocks of matter.
q Dalton’s model represented the atom as a featureless ball of uniform density.
q This model is referred to as the “soccer ball” model.
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DALTON’S���ATOMIC THEORY���
Dalton’s atomic theory, Ø explains the difference between an element
and a compound. Ø explains two scientific laws, and Ø predicts a new scientific law.
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DALTON’S ���ATOMIC THEORY���
Postulate Deduction Each element consists of indivisible, small particles called atoms.
All the atoms of an element are identical to one another, but different from others.
Gives a more precise definition for an element.
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5.1
Atoms consists of indivisible, small particles. Atoms of each element are identical to one another, but different from others.
DALTON’S���ATOMIC THEORY���
All atoms of oxygen are identical to one another
All atoms of hydrogen
are identical to one another
Atoms of oxygen are different
from atoms of hydrogen
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DALTON’S���ATOMIC THEORY���
Postulate Deduction Atoms combine chemically in definite whole-number ratios to form compounds.
Atoms can neither be created nor destroyed in chemical reactions.
Supports Law of Conservation of Mass.
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Supports Law of Definite Composition; predicts Law of Multiple Proportions.
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Atoms combine in definite whole-number ratios to form compounds.
H 2 = O 1
H 1 = O 1
LAW OF ���DEFINITE COMPOSITION���
As a result compounds always contain elements in the same proportions by mass.
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Two or more elements may combine in different ratios to form more than one compound.
H 2 = O 1
H 1 = O 1
LAW OF ���MULTIPLE PROPORTIONS���
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DISCOVERY OF���THE ELECTRON ���
q Smaller particles than the atom also exist and are called subatomic particles.
q In 1897, J.J. Thomson performed experiments with a cathode ray tube.
q Negatively charged particles from cathode were pulled towards positively charged plate, anode, and allowed to pass through and be detected on a fluorescent screen.
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DISCOVERY OF���THE ELECTRON ���
q In absence of a magnetic field, the cathode rays were not deflected.
q In presence of a magnetic and electric fields, the cathode rays were deflected towards the positive plate.
q These observations indicated that the cathode rays were negatively charged.
q These rays were later named electrons.
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ATOMIC���MODEL ���
q Based on these findings, Thomson proposed an atomic model, composed of negatively charged electrons embedded in a uniform positively charged sphere.
q This model is called the “plum pudding” model.
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DISCOVERY OF���THE NUCLEUS���
q In 1910, Ernest Rutherford carried out a number of experiments to further probe the nature of the atom.
q In these experiments he bombarded a thin sheet of gold foil with α-particles (large, positively charged) emitted from a radioactive source.
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DISCOVERY OF���THE NUCLEUS���
q The majority of the particles were observed to pass through un-deflected or slightly deflected.
q Some of the particles were observed to be deflected at large angles.
q Few of the particles were observed to be turned back towards the direction they came from.
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Deflection
Scattering
NUCLEAR MODEL���OF THE ATOM���
q Based on these observations, Rutherford proposed a model of the atom consisting of a small, massive positive center (nucleus), surrounded by electrons in mostly empty space.
q The deflections were caused by head-on collision of α-particles with the nucleus.
q The scatterings were caused by glancing collision of α-particles with the nucleus.
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THE MODERN��� ATOM���
q The current model of the atom describes it as a neutral spherical entity, composed of a positively charged nucleus surrounded by negatively charged electrons.
q The electrons (e-) move rapidly through the atomic volume, held by the attractive forces to the nucleus. q The nucleus consists of positively charged protons (p+) and neutrally charged neutrons (n0).
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ATOMIC���STRUCTURE���
q The modern atom consists of 3 subatomic particles:
Particle Charge Relative
Mass Proton +1 ~1800
Neutron 0 ~1800
Electron –1 1
q The number of protons in an atom determines its identity, and is called atomic number (Z).
q In a neutral atom, the number of protons (+) are equal to the number of electrons (–).
q Almost all the mass of the atom rests in the nucleus.
q Therefore the number of protons and neutrons in an atom is called the mass number (A).
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ATOMIC���STRUCTURE���
q The general designation for an atom is shown below:
Atomic number (Z) = # of protons
Mass number (A) = # of p+ + # of n0
# of n0 = A - Z
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ISOTOPES���
q Atoms of the same element that possess a different number of neutrons are called isotopes.
q Isotopes of an element have the same atomic number (Z), but a different mass number (A).
The 3 isotopes of Hydrogen
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ISOTOPES &���ATOMIC MASS���
q The mass of an atom is measured relative to the mass of a chosen standard (carbon-12 atom), and is expressed in atomic mass units (amu).
q The average atomic mass of an element is the mass of that element’s natural occurring isotopes weighted according to their abundance.
q Therefore the atomic mass of an element is closest to the mass of its most abundant isotope.
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Example 1:���
Determine the number of protons, electrons and neutrons in a chlorine atom .
3517Cl
A = 35Z = 17
# of protons = 17 (Z)
# of electrons = 17 (= p+)
# of neutrons = 18 (35 - 17)
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Example 2:���
Which two of the following are isotopes of each other?
410 410 412 412186 185 183 185X Y Z R
Isotopes of an element have the same atomic number, but a different mass number
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Example 3:���
Based on the information below, which is the most abundant isotope of boron (atomic mass = 10.8 amu) ?
Atomic mass of an element is closer to the mass of the more abundant isotope
Isotope 10B 11B
Mass (amu) 10.0 11.0
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(106.91) (0.5184) = 55.42 amu
(108.90) (0.4816) = 52.45 amu 107.87 amu
Isotope Mass (amu)
Abundance (%)
107Ag 106.91 51.84 109Ag 108.90 48.16
CALCULATING MASS���FROM ISOTOPIC DATA���
Atomic mass Abundance Mass of Abundance Mass ofx + x
of an element of isotope 1 isotope 1 of isotope 2 isotope 2
Ï ¸ Ï ¸Ï ¸ Ê ˆ Ê ˆ Ê ˆ Ê ˆÔ Ô Ô ÔÔ ÔÔ Ô Ô Ô Ô Ô˜ ˜ ˜ ˜Á Á Á Á= ˜ ˜ ˜ ˜Ì ˝ Ì ˝ Ì ˝Á Á Á Á˜ ˜ ˜ ˜˜ ˜ ˜ ˜Á Á Á ÁÔ Ô Ô Ô Ô ÔË ¯ˉ Ë ¯ˉ Ë ¯ˉ Ë ¯ˉÔ ÔÓ ˛ Ô Ô Ô ÔÓ ˛ Ó ˛
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THE END
