1. Chapter 2 Atoms

2. Classification of Matter
Figure 2.1 Classification of matter.

3. Classification of Matter
Compound: A pure substance that is made up of two or more elements in a fixed ratio by mass.
Formula of a compound: tells us the ratios of its constituent elements and identifies each element by its atomic symbol.
NaCl: the ratio of sodium atoms to chlorine atoms in sodium chloride is 1:1
H2O: the ratio of hydrogen atoms to oxygen atoms in water is 2:1

4. Classification of Matter
Element: A substance (for example, carbon, hydrogen, and iron) that consists of identical atoms.
There are 116 known elements.
Of these, 88 occur in nature; the others have been made by chemists and physicists.
Their symbols consist of one or two letters.
Names are derived from a variety of sources: the English name of the element, people important in atomic science, geographic locations, planets, mythological sources, etc.

5. A Water Molecule
Figure 2.2 Four representations of a water molecule.

6. Classification of Matter
Mixture: A combination of two or more pure substances.
The substances may be present in any mass ratio.
Each substance has a different set of physical properties.
Mixtures may be homogeneous or heterogeneous.
If we know the physical properties of the individual components of the mixture, we can use appropriate. physical means to separate the mixture into its component parts.

7. Dalton’s Atomic Theory
John Dalton ( )
All matter is composed of very tiny particles, which Dalton called atoms.
All atoms of the same element have the same chemical properties. Atoms of different elements have different chemical properties.
Compounds are formed by the chemical combination of two or more of the same or different kinds of atoms.
A molecule is a tightly bound combination of two or more atoms that acts as a single unit.

8. Evidence for Dalton’s Theory
Law of Conservation of Mass
Matter can be neither created nor destroyed.
As Dalton explained, if matter is made up of indestructible atoms, then any chemical reaction just changes the attachments among atoms, but does not destroy the atoms themselves.

9. Evidence for Dalton’s Theory
Law of Conservation of Mass
Monatomic elements consist of single atoms; for example, helium (He) and neon (Ne).
Diatomic elements: There are seven elements that occur as diatomic molecules:
H2, N2, O2, F2, Cl2, Br2, and I2
Polyatomic elements: Some elements have three or more atoms per molecule:
O3, P4, S8
Diamond has millions of carbon atoms bonded together to form one gigantic cluster.

10. Subatomic Particles The unit of mass is the atomic mass unit (amu).
One amu is defined as the mass of an atom of carbon with 6 protons and 6 neutrons in its nucleus.
1 amu = x g

11. A Typical Atom
Figure 2.6 Relative sizes of the atomic nucleus and an atom.

12. Mass and Atomic Number
Mass number: The sum of the number of protons and neutrons in the nucleus of an atom.
The mass of the electrons in an atom is so small compared to that of its protons and neutrons that electrons are not counted in determining mass number.
Atomic number: The number of protons in the nucleus of an atom.
A carbon atom of this composition is referred to as carbon-12.

13. Periodic Table
Figure 2.8 Four halogens

14. Isotopes
Isotopes: Atoms with the same number of protons but a different number of neutrons.
Carbon-12 has 6 protons and 6 neutrons
Carbon-13 has 6 protons and 7 neutrons
Carbon-14 has 6 protons and 8 neutrons
Most elements found on Earth are mixtures of isotopes.
Chlorine is 75.77% chlorine-35 (18 neutrons) and 24.23% chlorine-37 (20 neutrons).

15. Atomic Weight
Atomic weight: The weighted average of the masses (in amu) of the naturally occurring isotopes of an element.
Example: Chlorine is 75.77% chlorine-35 and 24.23% chlorine-37

16. Mass and Size of an Atom Consider an atom of lead-208.
It has 82 protons, 82 electrons, and 126 neutrons.
It has a mass of 3.5 x g.
It requires 1.3 x 1024 atoms to make 1 lb of lead-208.
The diameter of the nucleus is about 1.6 x m.
The diameter of the atom is 3.5 x m.
The density of the atom is 11.3 g/cm3.
The density of the nucleus is 1.8 x 1014 g/cm3.

17. 2 Periodic Table Dmitri Mendeleyev (1834-1907)
Arranged the known elements in order of increasing atomic weight beginning with hydrogen.
He observed that when elements are arranged in this manner, certain sets of properties recur periodically.
He then arranged elements with recurring sets of properties in the same column (vertical row); Li, Na, and K, for example, fall in the same column and start new periods (horizontal rows).

18. Figure 2. 8 The four halogens
Figure 2.8 The four halogens. Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.

19. Classification of the Elements
Figure 2.9 Classification of the elements.

20. Classification of the Elements
Metals
Are solids at room temperature (except for Hg, which is a liquid), shiny, conduct electricity, and are ductile and malleable.
Form alloys (solutions of one metal dissolved in another); brass, for example, is an alloy of copper and zinc.
In chemical reactions, they tend to give up electrons.
Nonmetals
Except for hydrogen (H), they lie on the right side of the Periodic Table.
Except for graphite, do not conduct electricity.
In chemical reactions, they tend to accept electrons.

21. Classification of the Elements
Metalloids
They have some of the properties of metals and some of nonmetals; for example, they are shiny like metals but, unlike metals, do not conduct electricity.
Six elements are classified as metalloids: boron, silicon, germanium, arsenic, antimony, and tellurium.
One of the metalloids, silicon, is a semiconductor; it does not conduct electricity under certain applied voltages, but becomes a conductor at higher applied voltages.

22. Examples of Periodicity

23. Examples of Periodicity

24. Examples of Periodicity

25. Electron Configuration
Electron configuration: The arrangement of electrons in the extranuclear space.
The energy of electrons in an atom is quantized, which means that an electron in an atom can have only certain allowed energies.
Ground-state electron configuration: The electron configuration of the lowest energy state of an atom.

26. Electron Configuration

27. Electron Configuration

28. Electron Configuration
Rule 2: Each orbital can hold up to two electrons with spins paired.
With four electrons, the 1s and 2s orbitals are filled and are written 1s2 2s2.
With an additional six electrons, the three 2p orbitals are filled and are written either 2px2 2py2 2pz2, or they may be written 2p6.

29. Electron Configuration
Figure 2.13 Energy levels for orbitals through the third shell.

30. Electron Configuration
Electron configurations are governed by three rules:
Rule 1: Orbitals fill in the order of increasing energy from lowest to highest.
Elements in the first, second, and third periods fill in the order 1s, 2s, 2p, 3s, and 3p.

31. Electron Configuration
Figure 3.12 The 1s, 2s, and 2p orbitals. Orbitals have definite shapes and orientations in space

32. Electron Configuration
Figure 2.14 The pairing of electron spins.

33. Electron Configuration
Rule 3: When there is a set of orbitals of equal energy, each orbital becomes half filled before any of them becomes completely filled.
Example: After the 1s and 2s orbitals are filled, a 5th electron is put into the 2px, a 6th into the 2py, and a 7th into the 2pz. Only after each 2p orbital has one electron is a second added to any 2p orbital.

34. Electron Configuration
Orbital box diagrams
A box represents an orbital.
An arrow represents an electron.
A pair of arrows with heads in opposite directions represents a pair of electrons with paired spins.
Example: carbon (atomic number 6)

35. Electron Configuration
Noble gas notation
The symbol of the noble gas immediately preceding the particular atom indicates the electron configuration of all filled shells
Example: carbon (atomic number 6)

36. Electron Configuration
Valence shell: The outermost incomplete shell.
Valence electron: An electron in the valence shell.
Lewis dot structure:
The symbol of the element represents the nucleus and filled shells.

37. Electron Configuration
Figure Electron configuration and the Periodic Table.

38. Electron Configuration

39. Periodic Property
As we have seen, the Periodic Table was constructed on the basis of trends (periodicity) in chemical properties.
With an understanding of electron configuration, chemists realized that the periodicity of chemical properties could be understood in terms of periodicity in electron configuration.
The Periodic Table worked because elements in the same column (group) have the same configuration in their outer shells.
We look at two periodic properties: Atomic size and ionization energy.

40. Atomic Size
The size of an atom is determined by the size of its outermost occupied orbital.
Example: The size of a chlorine atom is determined by the size of its three 3p orbitals, the size of a carbon atom is determined by the size of if its three 2p orbitals.

41. Atomic Size
Figure 2.16 Atomic radii of the main-group elements (in picometers).

42. Ionization Energy
Ionization energy: The energy required to remove the most loosely held electron from an atom in the gaseous state.
Example: When lithium loses one electron, it becomes a lithium ion; it still has three protons in its nucleus, but now only two electrons outside the nucleus, and therefore has a positive charge.

43. Ionization Energy
Figure 2.17 Ionization energy versus atomic energy for elements 1-37.

44. Ionization Energy Ionization energy is a periodic property:
In general, it increases across a row; valence electrons are in the same shell and subject to increasing attraction as the number of protons in the nucleus increases.
It increases going up a column; the valence electrons are in lower principle energy levels, which are closer to the nucleus and feel the nuclear charge more strongly.

45. Chapter 2 Atoms
End
Chapter 2