1. CHEM 100 Fall 2009 Chapter 2 Atoms, Molecules, and Ions

2. Atoms, Molecules, and Ions
CHEM Fall 2009
Atoms, Molecules, and Ions
The vocabulary of chemistry is specialized.
It includes such terms as atom, ion, molecule, isotope, acid, base, salt, and saturated hydrocarbon.
Chemistry also involves symbolic notation.
Atoms of elements are represented by symbols, such as H, C, N, O, Na, and Cl. These symbols are the alphabet of chemistry.
To represent compounds, symbols are combined into chemical formulas, such as NaCl, and formulas are the words of chemistry.
Chapter 2

3. Molecule models S 8 C 4 Gen. Chem. Chapter2 CHEM 100 Fall 2009

4. Molecular compounds Gen. Chem. Chapter2 CHEM 100 Fall 2009
Chemical formula – relative numbers of atoms of each element present
Empirical formula – the simplest whole number formula
Structural formula – the order and type of attachements
– shows multiple bonds
- may show lone pairs
- hard to show 3-d
Gen. Chem. Chapter2
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5. Some molecules H2O2 CH3CH2Cl P4O10 CH3-CH(OH)-CH3 HCOOH
CHEM Fall 2009
Some molecules
H2O2
CH3CH2Cl
P4O10
CH3-CH(OH)-CH3
HCOOH
Gen. Chem. Chapter2
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6. CHEM Fall 2009
Molecule models
Gen. Chem. Chapter2
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7. New laws are formulated making use of other laws.
CHEM Fall 2009
New laws are formulated making use of other laws.
BRICKS ON THE WALL!!!!
Gen. Chem. Chapter2
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8. more than the mass of the burning match?
Is the mass of the burned match pictured in the photograph the same as, less than, or
more than the mass of the burning match?
The match loses mass in the reaction. This kind of loss of mass to carbon dioxide during burning greatly confused early chemists who did not measure the mass of the gases produced.
The rule of conservation of matter is credited to Lavoisier.
Gen. Chem. Chapter2

9. Gen. Chem. Chapter2
Antoine Lavoisier (1743 – 1794)

10. CHEM Fall 2009
Another of the 18th century chemists to emphasize the importance of experiments was Antoine Lavoisier ( )
Lavoisier studying human respiration. His wife, Marie Anne Paulze, seated at the table on the right, records the experiment. Lavoisier was the first great theoretical chemist. Marie was an accomplished artist and made this drawing herself.
Gen. Chem. Chapter2
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11. 2.1 Laws of Chemical Combination
CHEM Fall 2009
2.1 Laws of Chemical Combination
Data obtained by experiment can often be summarized into laws. Scientific theories are then formulated to explain these laws.
Lavoisier: The Law of Conservation of Mass
Modern chemistry dates from the eighteenth century, when scientists began to make quantitative observations.
Antoine Lavoisier made mass measurements precise to about g and established that total mass does not change during a chemical reaction.
Gen. Chem. Chapter2
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12. law of conservation of mass.
CHEM Fall 2009
For example, Lavoisier heated the red oxide of mercury, causing it to decompose (break down) to two new products: mercury metal and oxygen gas.
By measuring carefully, he found that the total mass of the products was exactly the same as the mass of the mercury oxide he started with.
Lavoisier summarized his findings through the
law of conservation of mass.
The total mass remains constant during a chemical reaction.
Gen. Chem. Chapter2
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13. (also called the Law of Definite Proportions).
CHEM Fall 2009
Proust: The Law of Definite Proportions
By the end of the eighteenth century, Lavoisier and other scientists had succeeded in decomposing many compounds into the elements that form them.
One of these scientists, Joseph Proust (1754–1826), did careful quantitative studies by which he established the
Law of Constant Composition
(also called the Law of Definite Proportions).
All samples of a compound have the same composition; that is, all samples have the same proportions, by mass, of the elements present in the compound.
Gen. Chem. Chapter2
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14. Joseph Proust (1754 – 1826)
Gen. Chem. Chapter2

15. basic copper carbonate ;
CHEM Fall 2009
Proust based this law in part on his studies of a substance that we now call
basic copper carbonate ;
all samples of this compound have the same composition:
57.48% copper,
5.43% carbon,
0.91% hydrogen, and
36.18% oxygen by mass.
Gen. Chem. Chapter2
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16. CHEM Fall 2009
A substance known as basic copper carbonate occurs in nature
as the mineral malachite
forms as a patina on copper roofs and bronze statues
can be synthesized in the laboratory
The law of definite proportions : Regardless of its source, basic copper carbonate Cu2(CO3)(OH)2, always has the same composition.
Gen. Chem. Chapter2
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17. The Mineral Malachite
Chemistry: Cu2(CO3)(OH)2, Copper Carbonate Hydroxide.
Class: Carbonate
Gen. Chem. Chapter2

18. CHEM Fall 2009
A compound not only has a constant, or fixed, composition; it also has fixed properties.
For example, under normal atmospheric pressure, pure water always freezes at 0 °C and boils at 100 °C.
Gen. Chem. Chapter2
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19. CHEM Fall 2009
The physical and chemical properties of chemical substances—both elements and compounds—depend on their composition.
A twentieth-century modification of this law, based on the work of Albert Einstein, allows for the possibility of converting mass to energy (and vice versa).
This need not concern us, however, because in chemical rections the amounts of mass converted to energy are too small to detect.
Gen. Chem. Chapter2
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20. the Atomic Theory of Matter
CHEM Fall 2009
2.2 John Dalton and
the Atomic Theory of Matter
In 1803, John Dalton proposed a theory to explain
the laws of conservation of mass and constant composition.
As he developed what would become known as his atomic theory, Dalton found evidence of a scientific law describing the composition of matter.
In some cases, atoms of the same two elements are able to combine to form two or more different compounds, and the compositions of the different compounds are related.
Gen. Chem. Chapter2
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21. John Dalton
Gen. Chem. Chapter2

22. The Law of Multiple Proportions
CHEM Fall 2009
Let's illustrate Dalton's reasoning with a simple example: two compounds composed of the elements carbon and oxygen.
In carbon dioxide (the familiar gas produced in respiration and in the burning of fuels) the two elements combine in the ratio of 8.0 g of oxygen to 3.0 g of carbon.
In carbon monoxide (the poisonous gas formed when a fuel is burned in limited air) the elements combine in the ratio of 4.0 g of oxygen to 3.0 g of carbon.
Based on evidence such as this, Dalton formulated
The Law of Multiple Proportions
which we can state in the following way:
In two or more compounds of the same two elements, the masses of one element that combine with a fixed mass of the second element are in the ratio of small whole numbers.
Gen. Chem. Chapter2
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23. Law of Multiple Proportions
CHEM Fall 2009
Law of Multiple Proportions
Ratio of oxygen-to-carbon in CO2 is exactly twice the ratio in CO
Gen. Chem. Chapter2
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24. ANIMATION 1: MULTIPLE PROPORTIONS
Gen. Chem. Chapter2

25. Law of Multiple Proportions
CHEM Fall 2009
Law of Multiple Proportions
Four different oxides of nitrogen can be formed,
by combining 28 g of nitrogen with:
16 g oxygen, forming Compound I
48 g oxygen, forming Compound II
64 g oxygen, forming Compound III
80 g oxygen, forming Compound IV
What is the ratio 16:48:64:80 expressed as small whole numbers?
1 : 3 : 4 : 5
Compounds I–IV are N2O, N2O3, N2O4, N2O5
Gen. Chem. Chapter2
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26. CHEM Fall 2009
Dalton's Atomic Theory
The atomic model Dalton developed to explain the laws of chemical combination is based on four ideas:
- All matter is composed of extremely small, indivisible particles called atoms.
- All atoms of a given element are alike in mass and other properties, but the atoms of one element differ from the atoms of every other element.
- Compounds are formed when atoms of different elements unite in fixed proportions. That is, one atom of A to one of B in AB, two atoms of A to one of B in A2B and so on.
- A chemical reaction involves a rearrangement of atoms to produce new compounds. No atoms are created, destroyed, or broken apart in a chemical reaction.
Gen. Chem. Chapter2
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27. CHEM Fall 2009
2.3 The Divisible Atom
However, like most other scientific theories, Dalton's model eventually had to be modified in light of later discoveries.
Toward the end of the nineteenth century, certain experiments began to reveal that atoms are made up of smaller parts.
Of the dozens of subatomic (smaller than atomic) particles now known, three are of special importance in the study of chemistry.
They are the proton, neutron, and electron.
Gen. Chem. Chapter2
Chapter 2

28. CHEM Fall 2009
Subatomic Particles
The mass and charge of subatomic particles are so small that they are conveniently expressed in relative units.
The proton has a relative mass of 1. The proton also carries one fundamental unit of positive electric charge, denoted 1+ .
The neutron, as its name implies, is an electrically neutral particle; it has no charge. Although its mass is slightly greater than that of the proton, for many purposes we can consider the neutron also to have a relative mass of 1.
The third particle, the electron, has a mass that is 1/1836 of the mass of a proton. An electron has the same quantity of charge as a proton, but it is a negative charge, denoted as 1- .
Protons, neutrons, and electrons are fundamental particles.
This means that all protons are alike, all neutrons are alike, and all electrons are alike in whatever element they are found.
Gen. Chem. Chapter2
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29. CHEM Fall 2009
Subatomic Particles
Protons and neutrons are densely packed into a tiny, positively charged core of the atom known as the nucleus.
The extremely lightweight electrons are widely dispersed around the nucleus.
An atom as a whole is electrically neutral: It has no net charge because the negative and positive charges balance each other.
Atom like Football Field
Gen. Chem. Chapter2
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30. Nuclear Structure Atomic Diameter 10-8 cm Nuclear diameter 10-13 cm
CHEM Fall 2009
Nuclear Structure
Atomic Diameter cm
Nuclear diameter cm
= 1 Å
Gen. Chem. Chapter2
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31. CHEM Fall 2009
Subatomic Particles
Although the numbers differ from one element to another, the atoms of every element have equal numbers of electrons and protons.
Gen. Chem. Chapter2
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32. Every atom with 92 protons has Z = 92 and is an atom of uranium.
CHEM Fall 2009
Dalton believed that the identity of an element is determined by the mass of one of its atoms.
We now know that it is not the mass of an atom but rather the number of protons in the nucleus that determines the kind of atom and therefore the identity of an element.
The atomic number (Z) is the number of protons in the nucleus of an atom of a given element, and it is this number of protons that defines the element.
Every atom having two protons in its nucleus has Z = 2 and is an atom of helium.
Every atom with 92 protons has Z = 92 and is an atom of uranium.
Gen. Chem. Chapter2
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33. CHEM Fall 2009
Isotopes
Atoms that have the same number of protons but different numbers of neutrons are called isotopes.
The atomic number (Z) is the number of protons in the nucleus of a given atom of a given element.
The mass number (A) is an integral number that is the sum of the numbers of protons and neutrons in an atom.
The number of neutrons = A – Z.
Gen. Chem. Chapter2
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34. Isotopes For example, there are three isotopes of hydrogen.
CHEM Fall 2009
Isotopes
For example, there are three isotopes of hydrogen.
The most abundant isotope, occasionally called protium, has a single proton and no neutrons in its nucleus.
About one in every 6700 hydrogen atoms, however, has a neutron as well as a proton. Because the mass of a neutron is essentially the same as the mass of a proton, the mass of this hydrogen isotope, called deuterium, is about twice that of protium.
A third, very rare isotope of hydrogen, called tritium, has two neutrons and one proton in the nucleus. A tritium atom has about three times the mass of a protium atom.
Gen. Chem. Chapter2
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35. Isotopes of Hydrogen Atomic number (Z) = number of protons
CHEM Fall 2009
Isotopes of Hydrogen
Atomic number (Z) = number of protons
Protium has 1 proton, 0 neutrons : Z = 1
Deuterium has 1 proton, 1 neutron : Z = 1
Tritium has 1 proton, 2 neutrons : Z = 1
Gen. Chem. Chapter2
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36. CHEM Fall 2009
Some elements have only one naturally occurring isotope; these include fluorine-19, sodium-23, and phosphorus-31.
Most elements, however, have two or more isotopic forms.
Tin has the greatest number of naturally occurring isotopes: 10.
The naturally occurring isotopes are always found in certain precise proportions. In natural sources of chlorine, for example, 75.77% of the atoms are chlorine-35 and 24.23% are chlorine-37.
Chemical symbols for isotopes are commonly written in the form with A being the mass number and Z the atomic number of the element E.
For the two naturally occurring isotopes of chlorine, we can write and indicating mass numbers of 35 and 37
for chlorine-35 and chlorine-37, respectively. A E Z 35 Cl 17 37 Cl 17
Gen. Chem. Chapter2
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37. To represent a particular atom we use the symbolism:
CHEM Fall 2009
Chemical Symbol
To represent a particular atom we use the symbolism:
A= mass number Z = atomic number
Often do not specify Z when writing. For example 14C, C specifies Z = 12.
Special names for some isotopes. For example hydrogen, deuterium, tritium.
To represent an ion
Gen. Chem. Chapter2
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38. CHEM Fall 2009
Isotopes
Atoms can be represented using the element’s symbol and the mass number (A) and atomic number (Z): A E Z 35 Cl 17 37 Cl 17
How many protons are in chlorine-35?
How many protons are in chlorine-37?
How many neutrons are in chlorine-37?
Gen. Chem. Chapter2
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39. Other Examples of Isotopes
CHEM Fall 2009
Other Examples of Isotopes
The number of neutrons = A – Z
Carbon-12 Z = 6 so 6 neutrons
Carbon-13 Z = 6 so 7 neutrons
Carbon-14 Z = 6 so 8 neutrons
Chlorine-35 Z = 17 so 18 neutrons
Chlorine-37 Z = 17 so 20 neutrons
Uranium Z = 92 so 142 neutrons
Uranium Z = 92 so 143 neutrons
Uranium Z = 92 so 146 neutrons
Gen. Chem. Chapter2
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40. CHEM Fall 2009
C, H, O Analyser
Gen. Chem. Chapter2
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41. CHEM Fall 2009
2.4 Atomic Masses
By international agreement, the current atomic mass standard is the pure isotope carbon-12, which is assigned a mass of exactly 12 atomic mass units (12 u).
Based on this standard, we can define an atomic mass unit, (abbreviated amu and having the unit u)
as exactly one-twelfth the mass of a carbon-12 atom.
In more familiar units of mass, 1 u = x g.
Because we are interested in the masses of atoms and not their weights, we will now shift to the term atomic mass in place of the older atomic weight, except in a few cases of historical interest.
Gen. Chem. Chapter2
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42. CHEM Fall 2009
The atomic mass of an element is defined as the weighted average of the masses of the naturally occurring isotopes of that element.
Consider carbon, which consists of a mixture of two naturally occurring isotopes.
The much more abundant isotope is carbon-12. The other is carbon-13, with a mass of u. Both isotopes are present in substances containing carbon atoms, and in proportions that generally do not vary from one carbon-containing substance to another.
To describe the atomic mass of carbon, then, we need to use an average value, but not the simple average ( ) / 2 = Because carbon-12 is much more abundant than carbon-13, the average we seek lies much closer to the mass of carbon-12 than to that of carbon-13. We say that it is "weighted" toward the mass of carbon-12.
Gen. Chem. Chapter2
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43. To calculate the atomic mass of an element, we need two quantities—
CHEM Fall 2009
To calculate the atomic mass of an element, we need two quantities—
(1) the atomic masses of the isotopes of the element and
(2) the naturally occurring fractional abundances of the isotopes.
Let us explain how these quantities are obtained experimentally
To illustrate, let's return to the atomic mass of carbon.
Gen. Chem. Chapter2
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44. CHEM Fall 2009
The percentage abundances and fractional abundances of the carbon isotopes are as follows:
To obtain a weighted average atomic mass, we calculate the contribution of each isotope to the weighted average from the relationship
The term atomic weight is still widely used, however, as by the Commission on Atomic Weights of the International Union of Pure and Applied Chemistry (IUPAC).
Gen. Chem. Chapter2
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45. Contribution of carbon-12 = 0.98892 x 12.00000 u = 11.867 u
Use the data cited below to determine the weighted average atomic mass of carbon.
The contribution of each isotope to the weighted average atomic mass is given by Equation.
Solution
The contributions are
Contribution of carbon-12 = x u = u
Contribution of carbon-13 = x u = u
The weighted average atomic mass is the sum of the two contributions.
Atomic mass of carbon = u u = u
Gen. Chem. Chapter2

46. Carbon – 14 ( C-14 Dating )
Gen. Chem. Chapter2

47. 2.5 The Periodic Table: Elements Organized
CHEM Fall 2009
2.5 The Periodic Table: Elements Organized
In the nineteenth century, chemists discovered dozens of new elements.
By 1830, 55 elements were recognized, but there was no apparent pattern in their properties. Chemists badly needed a way to organize the growing collection of chemical data.
One way to do this was to arrange the elements in a manner that would establish categories of elements having similar physical and chemical properties.
Dimitri Mendeleev published the first successful arrangement, called a periodic table, in 1869.
In its modern form, the periodic table organizes a vast array of chemical knowledge.
Gen. Chem. Chapter2
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48. Dimitri Mendeleev
Gen. Chem. Chapter2

49. Mendeleev's Periodic Table Mendeleev arranged the elements in order of
CHEM Fall 2009
Mendeleev's Periodic Table
Mendeleev arranged the elements in order of
increasing atomic weight, from left to right in rows
and from top to bottom in columns (or groups).
In this arrangement, elements that most closely
resemble one another in physical and chemical
properties tend to fall in the same vertical group.
This group similarity repeats periodically, hence the name periodic table.
There would be no exceptions to the principle that all the elements in a group display similar properties, therefore Mendeleev placed some elements out of order, that is, not in the strict order of increasing atomic weight.
Gen. Chem. Chapter2
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50. For example, he correctly placed tellurium (atomic weight 127
For example, he correctly placed tellurium (atomic weight 127.6) before iodine (atomic weight 126.9) so that tellurium would be in the same column as the similar elements sulfur and selenium.
Gen. Chem. Chapter2

51. CHEM Fall 2009
When Mendeleev placed elements with similar properties in the same vertical group, a few gaps were created in his table.
Instead of seeing these gaps as defects, he boldly predicted the existence of undiscovered elements to fill the gaps.
Furthermore, because the table was based on patterns of properties, he was able to predict some properties of the missing elements.
Gen. Chem. Chapter2
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52. Mendeleev’s Table Mendeleevs’s early table was published in 1872.
CHEM Fall 2009
Mendeleev’s Table
Dimitri Mendeleev created this, the original periodic table.
Mendeleevs’s early table was published in 1872.
Using his table he was able to correct several values of atomic masses.
Because of its obvious usefullness his periodic table was almost universally adopted, and it remains one of the most valuable tools at the chemist’s use.
The only fundemantal difference between todays table and that of his is that in the current table the elements are ordered by atomic number rather than by atomic mass.
Stowe's table Spiral form Triangular form Folded Table Periodic Table
Gen. Chem. Chapter2
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53. CHEM Fall 2009
For example, he left a blank space for an undiscovered element that he called "eka-silicon" and used its location between silicon and tin to predict an atomic weight of 72 and other properties.
Gen. Chem. Chapter2
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54. CHEM Fall 2009
Table 2.2 shows just how accurate his predictions were, when compared to the properties of the actual element germanium discovered 15 years later.
The predictive nature of Mendeleev's periodic table led to its wide acceptance as a tremendous scientific accomplishment.
Gen. Chem. Chapter2
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55. The Periodic table Alkali Metals Noble Gases Alkaline Earths
CHEM Fall 2009
The Periodic table
Alkali Metals
Noble Gases
Alkaline Earths
Main Group
Halogens
Transition Metals
Lanthanides and Actinides
Gen. Chem. Chapter2
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56. (2) PERIODIC PROPERTIES ANIMATION
Gen. Chem. Chapter2

57. Empirical and Molecular Formulas
CHEM Fall 2009
Empirical and Molecular Formulas
Empirical formula: the simplest whole number ratio of elements in a compound
Example:
Molecular formula of glucose – C6H12O6
The elemental ratio C:H:O is 1:2:1, so the empirical formula is CH2O
Gen. Chem. Chapter2
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58. Molecular compounds Gen. Chem. Chapter2 CHEM 100 Fall 2009
Chemical formula – relative numbers of atoms of each element present
Empirical formula – the simplest whole number formula
Structural formula – the order and type of attachements
– shows multiple bonds
- may show lone pairs
- hard to show 3-d
Gen. Chem. Chapter2
Chapter 2

59. Introduction to Compounds
CHEM Fall 2009
Introduction to Compounds
A molecule is a group of two or more atoms held together by covalent bonds.
A chemical formula is a symbolic representation of the composition of a compound in terms of its constituent elements.
Gen. Chem. Chapter2
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60. Structural Formulas Shows how atoms are attached to one another.
CHEM Fall 2009
Structural Formulas
Shows how atoms are attached to one another.
Gen. Chem. Chapter2
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61. Molecular Compounds Ball-and-stick model vs. Space-filling model
CHEM Fall 2009
Molecular Compounds
Ball-and-stick model vs. Space-filling model
Gen. Chem. Chapter2
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62. Binary Molecular Compounds
CHEM Fall 2009
Binary Molecular Compounds
Molecular formulas are usually written with the more “metallic” first – “metallic” means farther left in the period and lower in the group
e.g., NaCl, KBr
Compounds that are typically comprised of two nonmetallic elements:
e.g., CO, NO, HF
Gen. Chem. Chapter2
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63. Formulas and Subscripts
CHEM Fall 2009
Formulas and Subscripts
Subscripts are used when a given atom is used more than once
e.g., H2O, CO2, N2O, HF, B2O3
The presence of subscripts is reflected in the names of compounds
Gen. Chem. Chapter2
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64. Names of Binary Compounds
CHEM Fall 2009
Names of Binary Compounds
Consider the compounds CO and CO2
The compound name consists of two words, one for each element in the compound
Name the element that appears first in the formula: CARBON
The second element has an altered name: retain the stem of the element name and replace the ending by -ide
OXYGEN
 OXIDE
However, both compounds cannot be carbon oxide
Gen. Chem. Chapter2
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65. Names of Binary Compounds
CHEM Fall 2009
Names of Binary Compounds
The names are further modified by adding prefixes to denote the numbers of atoms:
CO (Carbon Mon-oxide), CO2 (Carbon Di-oxide)
Gen. Chem. Chapter2
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66. CHEM Fall 2009
2.7 Ions and Ionic Compounds In an isolated atom, the number of protons equals the number of electrons, and the atom is therefore electrically neutral. In some chemical reactions, however, an individual atom or a group of bonded atoms may lose or gain one or more electrons, thereby acquiring a net electric charge and becoming an ion. Ions are formed only through the loss or gain of electrons; there is no change in the number of protons in the nucleus of the atom(s). If electrons are lost, there are more protons than electrons in the resulting ion and so it has a positive charge. If electrons are gained, there are more electrons than protons in the resulting ion and so it has a negative charge.
Gen. Chem. Chapter2
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67. Ions and Ionic Compounds
CHEM Fall 2009
Ions and Ionic Compounds
Atoms that gain or lose electrons are called ions
Positive ions: CATIONS Negative ions: ANIONS
Atoms that lose electrons form cations
Na  Na+ + e–
Atoms that gain electrons form anions
Cl + e–  Cl–
Gen. Chem. Chapter2
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68. CHEM Fall 2009
Monatomic Ions
Group A metals usually lose the number of electrons equal to their Group number.
Nonmetal atoms usually gain electrons and have a charge equal to their Group number minus eight.
The periodic table cannot be used to determine the charge on Group B metals.
For naming, Group B metals capable of multiple charges have the corresponding Roman numeral in parentheses added after the element name.
Gen. Chem. Chapter2
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69. CHEM Fall 2009
Common Monatomic Ions
FIGURE 2.10 Symbols and periodic table locations of some monatomic ions Three general observations can be made of these data:
Aluminum and the metals of groups 1A and 2A form just one cation, which carries a positive charge equal in magnitude to the A-group number.
Most of the metals of the B groups form two or more cations of different charges, though in some cases only one of these cations is commonly encountered.
The nonmetals of groups 7A and 6A, along with nitrogen and phosphorus of group A, form anions that have a charge equal to the group number minus 8.
Gen. Chem. Chapter2
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70. Formulas and Names for Ionic Compounds
CHEM Fall 2009
Formulas and Names for Ionic Compounds
Ionic compounds form when oppositely charged ions are attracted to each other
NaCl
Resulting compound is electrically neutral
Na+ Cl–
(+1) + (–1) = 0
Ionic compound names use the cation name followed by the anion name
Sodium chloride
Gen. Chem. Chapter2
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71. CHEM Fall 2009
Now consider the formula for aluminum oxide. We cannot simply combine one Al3+ and one O2- , because this would produce a formula unit with a net charge of The combination of two Al3+ ions and three O2- ions, though, is an electrically neutral formula unit: 2 (3+) + 3 (2-) = = 0 Therefore, the formula for aluminum oxide is Al2O3 .
Gen. Chem. Chapter2
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72. CHEM Fall 2009
Polyatomic Ions
Polyatomic ions are charged groups of covalently bonded atoms
Gen. Chem. Chapter2
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73. Hydrates Examples: BaCl2 . 2 H2O CuSO4 . 5 H2O
CHEM Fall 2009
Hydrates
A hydrate is an ionic compound in which the formula unit includes a fixed number of water molecules associated with cations and anions
Examples:
BaCl2 . 2 H2O
CuSO4 . 5 H2O
Gen. Chem. Chapter2
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74. Acids IntroToAcids Taste sour Turn blue litmus paper red
CHEM Fall 2009
Acids
Taste sour
Turn blue litmus paper red
React with metals to form hydrogen gas
Neutralize a base
IntroToAcids
Gen. Chem. Chapter2
Chapter 2

75. Bases IntroToBases Taste bitter Turn red litmus paper blue
CHEM Fall 2009
Bases
Taste bitter
Turn red litmus paper blue
Feel slippery on skin
Neutralize an acid
IntroToBases
Gen. Chem. Chapter2
Chapter 2

76. CHEM Fall 2009
Arrhenius Concepts
Acids are compounds that ionize in water to form a solution of H+ ions and anions
Bases are compounds that ionize in water to form solutions of OH– and cations
Gen. Chem. Chapter2
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77. Arrhenius Concepts Acids and bases react to form a salt and water
CHEM Fall 2009
Arrhenius Concepts
Acids and bases react to form a salt and water
= neutralization
HCl + NaOH  “Salt” + Water
Acid Base Na
cation Cl
/anion
HOH
Gen. Chem. Chapter2
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78. Animation (3) & (4) : Acids, Bases
Gen. Chem. Chapter2

79. Formulas and Names for Acids
CHEM Fall 2009
Formulas and Names for Acids
Binary acids start with hydro and end with “ic” plus the word acid
Ternary acids simply use the polyatomic anion name with “ate” changing to “ic” plus the word acid
Gen. Chem. Chapter2
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80. Formulas and Names for Bases
CHEM Fall 2009
Formulas and Names for Bases
Arrhenius bases always have hydroxide ions
The name follows ionic compound convention
e.g., NaOH – sodium hydroxide
Molecular bases form OH– after reacting with water
NH3 + HOH  NH4OH
Ammonia  ammonium hydroxide
Gen. Chem. Chapter2
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81. Formulas and Names for Salts
CHEM Fall 2009
Formulas and Names for Salts
Binary salts use the “ide” ending on the anion name
e.g., sodium chloride
Polyatomic salts use “ate” ending on the anion name
e.g., sodium sulfate
Gen. Chem. Chapter2
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82. Gen. Chem. Chapter2

83. CHEM Fall 2009
Organic Compounds
Organic Chemistry is the study of carbon and its compounds
Carbon compounds containing one or more of the elements H, O, N, or S are especially common
Most organic compounds are molecular compounds
Can exist as acids, bases, and salts
Compounds have systematic names AND common names
Gen. Chem. Chapter2
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84. Representations of Molecules
CHEM Fall 2009
Representations of Molecules
Condensed Structural Formula CH3CH2CH3
Structural Formula
Ball and Stick
Gen. Chem. Chapter2
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85. Saturated Hydrocarbons
CHEM Fall 2009
Saturated Hydrocarbons
Hydrocarbons have only hydrogen and carbon atoms
Saturated hydrocarbon: has
the maximum number of hydrogen
atoms possible for each carbon atom
Alkanes are saturated hydrocarbons
Methane (CH4) is the first molecule in the alkane series
Gen. Chem. Chapter2
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86. Prefixes for Number of Carbon
CHEM Fall 2009
Prefixes for Number of Carbon
Used for simple organic molecules
Combined with alkane ending “ane”
e.g., propane is a 3-carbon alkane
Gen. Chem. Chapter2
Chapter 2

87. CHEM Fall 2009
Isomers
Compounds with the same molecular formula but different structural formulas
Gen. Chem. Chapter2
Chapter 2

88. CHEM Fall 2009
FIGURE 2.14 Butane and isobutane Ball-and-stick and structural models illustrate the difference between the two isomers of butane.
Gen. Chem. Chapter2
Chapter 2

89. CHEM Fall 2009
Cyclic Alkanes
Alkane compounds that have carbons arranged in a ring structure are called cycloalkanes.
use the prefix cyclo-
cyclohexane
methylcyclopropane
Gen. Chem. Chapter2
Chapter 2

90. CHEM Fall 2009
FIGURE 2.15 Cyclohexane The molecular and line-angle formulas for cyclohexane indicate the bonding of atoms within the molecule. The ball-and-stick model further indicates that all of the carbon atoms of the cyclohexane molecule do not lie in the same plane. Rather, it can assume several different arrangements or conformations. The most stable arrangement, shown here, is called the chair conformation because the six carbon atoms outline a structure that somewhat resembles a reclining chair.
Gen. Chem. Chapter2
Chapter 2

91. CHEM Fall 2009
Functional Groups
Specific groupings of atoms attached to a carbon chain that give the compound unique properties
Most-common functional groups include:
Alcohols
Ethers
Carboxylic Acids
Esters
Amines
Gen. Chem. Chapter2
Chapter 2

92. Alcohols Alcohols are molecules that contain a hydroxyl group (OH)
CHEM Fall 2009
Alcohols
Alcohols are molecules that contain a hydroxyl group (OH)
Gen. Chem. Chapter2
Chapter 2

93. CHEM Fall 2009
Alcohols
Gen. Chem. Chapter2
Chapter 2

94. CHEM Fall 2009
Ethers
Ethers are molecules in which two alkane groups (R-) are attached to a central oxygen atom
The general formula is R-O-R´
CH3CH2OCH2CH3
CH3CH2OCH2CH2CH3
R and R´ may be the same or different groups
Gen. Chem. Chapter2
Chapter 2

95. Carboxylic Acids HCOO– + H+
CHEM Fall 2009
Carboxylic Acids
Carboxylic acids are alkanes that also contain a carboxyl group and are weak acids
Acts like an Arrhenius acid, loses a hydrogen ion
HCOO– + H+
Gen. Chem. Chapter2
Chapter 2

96. CHEM Fall 2009
Gen. Chem. Chapter2
Chapter 2

97. CHEM Fall 2009
Esters
Esters are molecules in which two alkanes are attached to each side of a carboxyl group
(R’-COO-R)
Gen. Chem. Chapter2
Chapter 2

98. CHEM Fall 2009
Amines
Amines are molecules in which alkanes and hydrogen(s) are attached to a central nitrogen
NH2(CH2)4NH2
Amines are weak bases
Gen. Chem. Chapter2
Chapter 2

99. CHEM Fall 2009
Summary of Concepts
The basic laws of chemical combination are the laws of conservation of mass, constant composition, and multiple proportions.
The three main subatomic particles are the protons, neutrons, and electrons.
Atoms with the same number of protons but different numbers of neutrons are called isotopes.
A chemical formula indicates the relative numbers of atoms of each type in a compound.
Gen. Chem. Chapter2
Chapter 2

100. CHEM Fall 2009
Summary (cont.)
The periodic table is an arrangement of the elements by atomic number that places elements with similar properties into the same vertical group.
Ions are formed by the gain or loss of electrons. Positive ions are cations and negative ions are anions.
Many compounds are classified as either acids (H+), bases (OH–), or salts (neutralization of acid and base).
Organic compounds are based on the element carbon.
Functional groups confer distinctive properties on an organic molecule.
Gen. Chem. Chapter2
Chapter 2

101. Arrhenius Concepts Acids and bases react to form a salt and water
CHEM Fall 2009
Arrhenius Concepts
Acids and bases react to form a salt and water
= neutralization
HCl + NaOH  “Salt” + Water
Acid Base Na
cation Cl
/anion
HOH
Gen. Chem. Chapter2
Chapter 2