Download File

Handout in General Chemistry
Subject matter: ATOMS, MOLECULES, IONS
Matter is made up of molecules. These are either molecules of an element or of a compound.
Compound is made up of two or more elements. In turn, molecules are made up of atom. An atom is
the smallest fundamental particle of an element that has the properties of that element. (on the
other hand, a molecule is formed by the chemical combination of two or more atoms.)
I.
II.
The Atomic Theory
1. Democritus (Greek philosopher) – (460 – 370 B.C.) and Leucippus (c. 500 B.C.) first
proposed that matter is made up of atoms.
- Atom comes from Greek word “atomos” which means “indivisible”
- Democritus coined the name “atom. He hypothesized from observations that :
a) Atoms were small, hard particles made of the same material but of different shapes
and sizes;
b) There were infinite number of these atoms and they are in const motion;
c) Atoms have the ability to combine with other atoms
d) Atoms could no longer be divided into smaller particles
2. John Dalton (English school teacher and amateur meteorologist), based on experimental
evidence, proposed in 1803 that:
a) Matter is composed of small, indivisible particles called atoms;
b) Atoms of the same elements are identical and have the same properties;
c) Chemical compounds are composed of atoms of different elements combined in small
whole-number ratios;
d) Chemical reactions are merely the rearrangement of atoms into different
combinations.
3. Joseph John Thomson (English physicist – 1856 -1940), after experimenting with charged
particles, proposed that:
a) The atom is not indivisible
b) The atom has an negatively charged particle which he called “corpuscles”
c) Inferred that there must also be positively charged particle in the atom
d) His atomic model is also called plum pudding model (more like an atis)
e) The atom is made of positively charged material (the flesh of the atis) Spread out in
this positively charged material were the negatively charged electrons (seeds of the
atis)
4. Sir Ernest Rutherford (English physicist) In 1910, proposed:
a) The atom consisted of a small, dense and positively charged center he called
“nucleus”
b) Distributed in the space surrounding the nucleus are the electrons.
- He found the positively charged proton in 1919. A neutral particle, the neutron, was
discovered in 1932 by another scientist.
5. Niels Bohr (Danish scientist – 1885 – 1962) In 1913, proposed that:
a) The electrons move around the nucleus much like the planets revolve around the sun
in orbits
b) For this reason, his model is called “planetary model” of the atom
c) Electron orbits are known as “energy levels’
6. Quantum Mechanical Model: ‘Clouding’ the electron
a) Atom has the nucleus at the center, made up of positively charged particle
‘proton’ and neutrally charged particle ‘neutron’;
b) Around the nucleus is the electron cloud with the electrons
c) The electrons are still located in energy levels, but instead of orbits, the
energy levels are the regions that make up the electron cloud.
The Structure of the Atom
- An atom has a nucleus at the center made up of positively charged ‘proton’ and neutral
‘neutron’. It has electron cloud surrounding the nucleus where the negatively charged
‘electron’ moves.
- Today , > 200 different subatomic particles have been discovered.
1
Handout in General Chemistry
- Orbital – The region of space where there is the significant probability of finding a particular
electron.
- S orbital – a spherical electron cloud; can accommodate 2 electrons
- P orbital – electron cloud shaped with two lobes on either side of the nucleus. There are three
p-orbitals: px py, pz Can accommodate 2 electrons per p-orbital.
- D orbital – electron cloud shaped with 4 lobes. There are 5 d-orbitals: dxy, dyz, dx2-y2, dz2 Can
accommodate 2 electrons per d-orbital
- F orbital – There are 7 f-orbitals. Can accommodate 2 electrons per f-orbital
- Subshell – orbitals of the same type. Ex. S –orbital is a subshell, p-orbitals are subshells
- Electron clouds are divided into energy level. The farther from the nucleus, the higher the
energy level.
- Each cloud level is called shell. Each shell has subshells.
- The 1st shell has only one subshell (1s) : s-orbital which has 2 electron capacity
- The 2nd shell has 2 subshells (2s, 2p) : s – orbital and p – orbital (with 3 shapes)
- The 3rd shell has 3 subshells (3s, 3p, 3d) : s-orbital, p-orbital, d-orbital (w/ 5 shapes)
- The 4th shell has 4 subshells (4s, 4p, 4d, 4f): s-orbital, p-orbital, d-orbital, f-orbital w/ 10
shapes)
- Each orbital can accommodate 2 electrons, thus the 1st shell has 2 electron capacity, the 2nd
shell has 8 electron capacity, the 3rd shell has 18 electron capacity, the 4th shell has 32.
- The total electron capacity of a shell is equal to 2n2 where n is the shell number.
III.
Atomic Number, Mass Number, Isotopes
- Size a typical atom would have a diameter of 10-8 cm.
- Mass of a proton is 1.67 x 10-24 g.
- The neutron’s mass is slightly greater but is considered equal to the proton’s mass.
A. Atomic Number
- Atomic number of an element is the number of protons of its atom.
- It is the atomic number that distinguishes one element from another. Any atom with an atomic
number of 29, regardless of any consideration, is an atom of copper. If the atomic number is
28, the element is nickel; if it is 30, the element is zinc.
- Ex. The simplest element, hydrogen, has an atomic number of 1. It has only one proton in the
nucleus. Since it is neutral, it also has only one electron. An atom with two protons in its
nucleus has an atomic number of 2. This is helium.
- Presently, there are 113 known elements with atomic numbers 93 – 112 and 114 are
artificially produced.
B. Mass Number
- Atomic mass unit (amu) – for convenience, scientists have agreed to establish a standard unit
of measure for atomic mass called ‘amu’. It is based on the mass of the most common
carbon form which is assigned a mass of 12 amu.
- Mass number is the sum of the number of protons and neutrons.
- This is usually used to distinguish the isotopes of an element from each other.
C. Isotopes
- Isotopes – atoms of the same element with different number of neutrons.
- Most elements in nature exist as a mixture of isotopes
- Ex. Hydrogen always has one proton in its nucleus but it has 3 different isotopes. Protium, the
most common, has no neutrons; Deuterium has one neutron; while tritium has two.
- The convention for verbally naming specific isotopes is to use the element’s name followed
by its mass number. Ex. copper-63 and copper-65. What does the number imply?
- Isotopic mass is the mass of the particular isotope of an element.
- Isotopic Notation of copper 6329Cu where 63 is the mass number and 29 is the atomic
number. Exercises: What is the atomic no. and mass no. of the ff: 2713Al; 168O; 1 1H; 126C
- Ex. Isotopes of an element are different in terms of their mass. How will they compare in
terms of their chemical properties?
D. Atomic mass
- In nature, a certain isotope of a given element will be more abundant than the others.
- Atomic mass – the relative abundance of the isotopes of a given element.
2
Handout in General Chemistry
- To get the atomic mass, multiply the mass of each isotope with its abundance, then get the
average.
- Ex. In nature, the element boron occurs as 19.9% 10B and 80.1% 11B. If the isotopic mass of
10
B is 10.013 amu and that of 11B is 11.009 amu. What is the atomic mass of boron?
E. Assessment
IV.
The Periodic Table
- Elements are materials made up of only one kind of atom.
- Elements are presented in the Periodic Table, arranged accdg to atomic no. (no. of protons)
- Periodic Law – the properties of elements are periodic functions of their atomic number.
- Earliest version was introduced in 1869 by Dmitri Mendeleev of Russia arranged according to
atomic mass
- Lothar Meyer of Germany independently presented a similar table in 1870.
- Henry Moseley, a British physicist, carried on the work of Mendeleev after his death and
arranged the elements according to atomic no. instead of atomic mass.
- Indicated in the present day table are the element symbol, atomic no. and atomic mass
- As of September 2003, 113 elements have been identified. The creation of element 114 was
announced at the end of 1998 but the claim has still to be ratified by the Int’l Union of Pure
Applied Chemistry (IUPAC)
- Only 99 different kinds of atoms make up all natural occurring substances. Fifteen more have
been artificially produced in laboratories.
A. Three (3) Main Categories
- Elements in the table are divided into 3 main categories: metals, nonmetals, metalloids.
- Metals are elements that have metallic properties like electrical & thermal conductivity, luster
- Metalloids are elements that have both metallic and non-metallic properties like silicon which
has no luster but can conduct electricity.
- The jagged line at the right side of the table separates these three categories of elements.
B. Periods
- Horizontal rows of elements in the periodic table are called periods
- Each period ends with a member of the family called noble gases.
- The atomic numbers of elements in a period increases from left to right.
C. Groups
- Families of elements falling into vertical columns are called groups or families.
- Each group is designated a number at the top of the group in Roman numeral which increases
from left to right. The Roman numeral is followed by the letter A or B
- Members of a group have similar properties.
- The atomic numbers in a group increases from top to bottom
- The groups of elements can be classified further into 4 categories: (Group categories)
1. The main Group or Representative Elements (Groups IA – VIIA)
a) Group IA - alkali metals (Li, Na, K excluding hydrogen which is a non-metal)
- Highly reactive metals meaning they combine easily with other elements.
- Have one electron in the highest energy level of their atoms thus it is easily lost.
Consequently, alkali metals are found in nature as positively-charged ions to form salts.
- In pure form are so soft that they can be cut easily with a knife. They have low densities
and melt at low temperatures.
b) Group IIA – Alkaline Earth Metals (Ca, Mg are most common)
- Reactive but not as reactive as elements in Group IA.
- Have two electrons in the highest energy level.
- Occur in nature as ions combined with other elements.
- Used to make building materials.
c) Group VIIA – Halogens (F, Cl, Br,) They are all nonmetals.
d) Group IIIA, IVA, VA – Generally, not referred to by a family name but by the element
at the top of the column. Ex. Elements in VA is part of the nitrogen family.
2. The Noble or Inert Gases (Group VIIIA)
- These elements form few chemical compounds. He, Ne, Ar do not form any compound.
3. The Transition Metals (Group B Elements)
3
Handout in General Chemistry
-
These elements are hard and shiny, with high melting points and are good conductors of
heat and electricity.
- Transition metals can combine with other elements but are less reactive than Alkali and
Alkaline Earth Metals.
- Transition metals are often found in ores. Ores are minerals containing relatively large
amounts of metal compounds.
- Transition metals make up most metal objects. Some form colorful compounds.
4.The Inner Transition Metals
- the 14 inner transition metals between lanthanium (atomic no. 57) and hafnium (atomic
noo. 72) are known as the lanthanides or rare earth, and the 14 metals between
actinium (atomic no. 89) and rutherfordium (atomic no. 104) are known as actinides.
- Except for element 61, all of the lanthanides occur in nature. They are called “rareearths” because they occur only in small amounts in the earth’s crust.
- Most of the actinides are synthetic, or made in the laboratories.
D. Physical States and the Periodic Table
- At room temperature (25°C) All 3 phases of matter are found among the elements in the
periodic table. Gas : all noble gases + hydrogen, nitrogen, oxygen, fluorine, chlorine
Liquid : mercury (a metal) and bromine (a nonmetal). All other elements are solids
- All the gaseous elements except for the noble gases which exist as individual atoms, are
composed of diatomic molecules (e.g. N2, O2, and H2)
- Halogens are diatomic. (F2, Cl2, Br2, I2)
E. Assessment
V. Molecules and Ions
A. Molecules
- Molecules are made up of two or more atoms. Ex. Oxygen molecule – O2 or hydrogen
molecule – H2, water molecule – H2O. Note: There is no mention on the kind of atoms.
- Compounds – Pure substances made up of molecules. It is made of two or more elements,
chemically combined.
- Elements vs Compounds : Elements have only one kind of atom while compounds have
two or more kinds of atoms. Note: it does not talk of no. of atoms but rather kinds of
atoms. Thus, elements can have more than one atom for as long as they are all of the same
kind. H2, Cu are examples of elements.
- Ex. State whether an element or a compound: O2, C, CO2, NaCl, P4 , FeCl3
- Atoms vs Molecules : Molecules are made up of two or more atoms; they can be of the
same kind or different. Note: This time, it talks of the no. of atoms. H2 and H2O are
examples of molecules, Cu is an example of an atom.
- Ex. State whether a molecule or an atom: O2, C, CO2, NaCl, S, P4 , FeCl3, Au, Hg
- Most substances in nature occur as molecules, not necessarily compounds. Take the case
of air which is made up of molecules of Oxygen, Nitrogen, Hydrogen, etc.
- Ex. State whether the statement is true or false: a) All compounds are made up of
molecules, b) All molecules are compounds
B. Ions
- Ions – a charged particle formed when an atom or a group of atoms gains or loses one or
more electrons. A positive ion has lost one or more electrons, and a negative ion has
gained one or more electrons.
- An ionic compound is formed by the combination of positive and negative ions. The ions
are held together by electrical attraction.
- Some elements have the tendency to lose electrons because of the few electrons in their
highest energy levels or shell (orbits)
C. Electron Configuration
- The electrons of an element are arranged in its energy levels starting with the lowest level
found closest to the nucleus upto the highest level that is farthest from the nucleus.
- A max. of two electron can be accommodated in the 1st shell, 8 in the 2nd shell, 18 in the
3rd shell and 32 in the 4th shell, etc.
- Aufbau Principle – electrons occupy the available orbitals in the subshells of lowest
energy.
4
Handout in General Chemistry
Electron configuration – the assignment of all the electrons in an atom into specific shells
and subshells.
- A shell is designated by the principal quantum number (n) and the subshell by the
appropriate letter. The number of electrons in that subshell can be displayed with an
appropriate superscript number. Ex. 4p3 where 4 refers to the shell, p refers to the
subshell (p-orbital) and 3 refers to the number of electrons in the subshell.
H: 1s1, He: 1s2, C: 1s22s22p2, Kr: 4s23d104p6
- Electron Configurations of Specific Groups
Fill up the orbitals in the ff order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f
In general: Representative Elements – [NG] nsxnpy
Noble Gases – [NG] ns2np6
Transition Metals – [NG] ns2(n-1)dx
Inner Transition Metals – [NG] ns2(n-1)d1(n-2)fx
D. Chemical Bonding
- When atoms combine chemically, they create a chemical bond. A chemical bond is an
attractive force.
- An atom with a full outer energy level is more stable than an atom with an unfilled outer
level. Atoms form bonds to become more stable.
- While protons and neutrons make up most of the atom’s mass, they contribute little to an
element’s chemical properties. It is the electrons that determine an element’s chemical
behaviour.
- It is the arrangement of electrons in the outermost energy level of an atom that dictates
how it bonds with other atoms. Since the electrons in the outermost energy level (shell)
are farthest from the nucleus, there is less attraction between these electrons and the
protons. Thus, these electrons can interact with other atoms.
- Octet Rule – Elements of the representative elements form bonds so as to have access
to eight outer electrons.
- Valence electrons – the outer s and p electrons in the atoms of a representative element.
Only the valence electrons are involved in bonding.
- Lewis dot symbols represent valence electrons as dots around the symbol of the element.
- The 1st four valence electrons are placed around the symbol, one dot at a time, on each
side of the four sides. The next four valence electron dots are then paired up with each of
the originals.
- Elements in each group have the same number of valence electrons and thus the same
number of dots.The Roman numeral of the group number also represents the number of
valence electrons.
- Ex. Write the Lewis dot symbol of 2Be, 8O, 7N, 6C, 9F, 15P, 16S, 17Cl
- Ways to form a stable octet configuration
a) a metal may lose 1 – 3 electrons to form a cation with the electron configuration of
the previous noble gas.
b) a nonmetal may gain 1 – 3 electrons to form an anion with the electron configuration
of the next noble gas.
c) Atoms (usually two nonmetals) may share electrons with other atoms to obtain access
to the number of electrons in the next noble gas.
- There are three different kinds of chemical bonds.
a) Ionic bonds – electrons are transferred from atom to atom; these are electrostatic
attractions between oppositely charged ions. Usually formed between metals and
nonmetals. Ionic compounds are nearly always solids.
b) Covalent bonds – electrons are shared. Usually formed between the same nonmetals.
Covalently bonded compounds are more common than ionic compounds.
c) Metallic bonds – involve the sharing of electrons among many atoms
Certain molecules held together by covalent bonds tend to gain or lose electrons as a unit.
The whole group of atoms becomes an ion with a positive or negative charge. These are
called polyatomic ions.
-
VI.
Chemical Formulas
5
Handout in General Chemistry
-
-
A chemical formula is a combination of symbols and numbers that represent the
composition of a compound.
The symbols refer to the kinds of atoms in the compound.
The number, called subscript, refers to the number of each kind of atom.
When more than one atom of an element is present in a compound, a subscript is written
to the right and below the element’s symbol. If there is only one atom of the element, no
subscript is used.
Remember: A compound consist of more than one element.
Ex. Water – H2O, Carbon dioxide- CO2, Methane- CH4 , Ammonia – NH3
VII. Naming Compounds / Ions
- If a metal cation is named alone, the word ‘ion’ is added Ex. Na+ - sodium ion, H+
hydrogen ion.
- In naming and writing the formula for binary ionic compounds, give the English name of
the metal cation first, then give the name of the anion (English root + ‘ide’) Ex. NaCl –
sodium chloride, CO2 – Carbon dioxide.
- Stock method of naming: There are cases when a metal forms more than one cation, thus
more than one compound is formed. To distinguish between the two compounds, the
charge on the metal ion follows the name of the metal in Roman numerals and in
parenthesis. Ex. iron (II) chloride – refers to FeCl2, iron (III) chloride – refers to FeCl3
- Classical method of naming: No longer widely used in chemistry. The name of the metal
ion that has the lower charge ends in ‘-ous’ and the higher ends in ‘-ic’.
- If the symbol of the element is derived from a Latin word, the Latin root is generally used
rather than the English root. Ex. ferrous chloride and ferric chloride
Monoatomic Ions of the Representative Elements
Ion
Name
Ion
Name
Ion
+
2+
H
Hydride
Be
Beryllium
C4+
2+
Li
Lithium
Mg
Magnesium
N3+
2+
Na
Sodium
Ca
Calcium
P3+
2+
K
Potassium
Sr
Strontium
O2Rb+ Rubidium
Ba2+
Barium
S2+
3+
Ce
Cesium
Al
Aluminum
Se2Metals that Form More Than One Ion
Ion
Stock Name
Classical Name
Cr2+
chromium (II)
chromous
Cr3+
chromium (III)
chromic
2+
Fe
iron (II)
ferrous
Fe3+
iron (III)
ferric
Pb2+
lead (II)
plumbous
Pb4+
lead (IV)
plumbic
+
Au
gold (I)
aurous
Au3+ gold (III)
auric
Polyatomic Ions
Ion
Name
C2H3O2- Acetate
NH4+
Ammonium
2CO3
Carbonate
ClO4Perchlorate
ClO3Chlorate
ClO2Chlorite
ClO
hypochlorite
CrO42Chromate
Ion
Cr2O72CNH2PO4HCO3HPO42HSO4HSO3-
Ion
Cu+
Cu2+
Mn2+
Mn3+
Sn2+
Sn3+
Co2+
Co3+
Name
Carbide
Nitride
Phosphide
Oxide
Sulfide
Selenide
Stock Name
copper (I)
copper (II)
manganese (II)
manganese (III)
tin (II)
tin (III)
cobalt (II)
cobalt (III)
Name
Dichromate
Cyanide
Dihydrogen phosphate
Hydrogen carbonate/
Bicarbonate
Monohydrogen phosphate
Hydrogen sulfate / bisulfate
Hydrogen sulfite / bisulfite
Ion
OHNO3NO2C2O42MnO4PO43SO42SO32-
Ion
Te2FClBrI-
Name
Telluride
Fluoride
Chloride
Bromide
Iodide
Classical Name
cuprous
cupric
manganous
manganic
stannous
stannic
cobaltous
cobaltic
Name
Hydroxide
Nitrate
Nitrite
Oxalate
Permanganate
Phosphate
Sulfate
Sulfite
6
Handout in General Chemistry
Greek Prefixes:
Number
1
2
3
4
5
Prefix
monoditritetrapenta-
Number
6
7
8
9
10
Prefix
hexaheptaoctanonadeca-
Exercises:
1. Name the ff compounds: a) Li2O b) CrI3 c) PbS d) Mg3N2 e) Ni3P2 f) HC2H3O2
g) HNO2 h) HI
2. Provide the formula for the ff compounds: a) aluminium iodide b) iron (III) oxide
c) tin (IV) bromide d) calcium nitride e) barium hydroxide f) silver (I) hypochlorite
Reference Material :
Malone, Leo, and Theodore Dolter. Basic Concepts of Chemistry. 8th ed.
7
Handout in General Chemistry
III. Atomic Number, Mass Number and Isotopes Assessment
1. Describe the present day concept of an atom’s structure.
2. Identify the following as atomic number, mass number, or isotopic mass:
a) always an integer value
b) the superscript in isotopic notation,
c) the total mass of the atom
d) the number of protons
e) determines the number of protons and neutrons combined
3. Write the isotopic notation for the following:
a) element X : protons=9, neutrons=10, electrons=9
b) element Y : protons=35, neutrons=44, electrons=35
c) element Z : protons=20, neutrons=20, electrons=20
4. The most common isotope of carbon has 6 protons and 6 neutrons, giving a total mass of
12 amu. In terms of amu, what is the mass of a proton? A neutron?
5. Uranium has an atomic number of 92. Two of its isotopes are uranium-235 and uranium238. How many neutrons are in the nuclei of each of these isotopes?
6. The atomic mass of carbon is 12.011 amu. Which of its isotopes, carbon-12 or carbon-14,
is more abundant in nature? Explain.
7. Lithium has two naturally occurring isotopes: 6Li has a mass of 6.015 amu and is 7.42%
abundant; 7Li has a mass of 7.016 amu and is 92.58% abundant. Calculate the atomic
mass of lithium.
IV. Periodic Table Assessment
2. What is the most common physical state of the elements in the periodic table at room
temperature?
3. Which are more common, metals or nonmetals?
4. Why do alkali metals and alkaline earth metals exist in nature as ions?
5. Where might you find calcium in your school?
6. Of the three main categories of elements, to what category does nickel (Ni) belong?
Sulfur (S)? Germanium (Ge)?
7. In the periodic table, elements are arranged according to increasing ________________.
8. An element in Group IIA, such as calcium (Ca), is also known as __________________.
9. An element in Group IA, such as Lithium (Li), is also known as _________________.
10. Assign each of the following elements to one of the four group categories, give its
physical state, tell whether a metal or nonmetal, give its period and its group number
a) copper (Cu) b) Argon (Ar) c) Barium (Ba) d) Bromine (Br) e) Uranium (U)
11. Identify the element that fits each description
a) fourth period alkaline earth metal
b) Liquid halogen
c) IIIA metalloid
d) The last transition metal of the fifth period
e) The first period noble gas.
VI/VII Chemical Formulas / Naming of Compounds
1. What are the chemical formula of the ff compounds and state the naming method used :
a) nickel (II) cyanide, b) aluminum sulfite, c) magnesium bicarbonate, d) iron (III) perchlorate,
e) ammonium phosphate
2. Name the following compounds: a) K2CO3,
b) Fe2(SO4)3 c) Ba(OH)2
d) LiH2PO4
e) AgClO, f) K2CrO4
g) Co2(CO3)3
8