Chapter 2 Atoms, Ions, and the Periodic Table notes/Microsoft... · Mendeleev’s Table • Russian chemist Dmitri Mendeleev developed and published the basic arrangement of the periodic

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Chapter 2Atoms, Ions, and the Periodic Table

• Dalton’s Atomic Theory

• Structure of the Atom

• Ions

• Atomic Mass

• The Periodic Table

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Law of Conservation of Mass

• the mass of the products (overall)

always equals the mass of the

reacting substances

• Proposed by Antoine Lavoisier in

1787

• His experiments showed that no

measurable change in mass occurs

during a chemical reaction.

2 -How could you show mass is conserved in a reaction?

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Law of Conservation of Mass

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Law of Definite Proportions

• Proposed by Joseph Proust between 1797 and 1804

• states that all samples of the same

compound always contain the same proportions by mass of the component

elements

– For example, water is always composed of oxygen and hydrogen in a mass ratio of 8:1 (or 16:2).

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Dalton’s Atomic Theory

• Dalton’s Atomic Theory has 4 postulates:1. All matter is composed of exceedingly small,

indivisible particles called atoms.

2. All atoms of a given element are identical both in mass and in chemical properties. However, atoms of different elements have different masses and different chemical properties.

3. Atoms are not created or destroyed in chemical reactions.

4. Atoms combine in simple, fixed, whole-number ratios to form compounds.

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Structure of the Atom• Atoms actually are divisible. They are composed of

subatomic particles.

• Subatomic particles include:

– 1 kind of particle found outside the nucleus

• Electrons

– negatively charged subatomic particles

– 2 kinds of particles found in the nucleus (center of the atom)

• Protons

– positively charged subatomic particles

• Neutrons

– uncharged subatomic particles2 -

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Structure of the Atom

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Figure 2.10

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The Discovery of Electrons• The existence of the electron was demonstrated by

J.J. Thomson in 1897.– He conducted a series of experiments with cathode ray

tubes, in which:

1. Voltage was applied by connecting each end of a tube to a battery.

2. The electricity forms rays that flow from one end of the tube to the other and that are visible through specially coated glass.

3. When an electric or magnetic field was applied to the tube (and the rays), the rays bent toward a positively charged plate, and were deflected by a negatively charged plate. Because like charges repel and opposite charges attract, the particles were negatively charged.

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The Discovery of Electrons

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Figure 2.6

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The Nuclear Atom• From his experiments with electrons, J.J. Thomson proposed

that electrons might be embedded in a sphere of positive charge (“plum pudding” model of the atom).

• Ernest Rutherford designed an experiment in the early 1900’s to test J.J. Thomson’s “plum pudding” model of the atom.

– The experiment involved bombarding a piece of gold foil with alpha particles (positively charged Helium atoms without the electrons).

– Alpha particles were expected to zip through the gold foil, and most did, but some were deflected slightly and a few bounced backwards.

– The deflected particles led to the hypothesis of the nucleus, a concentrated, positively charged core, while electrons occupied the volume outside of the nucleus.

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The Nuclear Atom

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Subatomic Particles• Protons have:

– a charge equal to +1.6022 x 10-19 C (expressed as +1)

– a mass equal to 1.6726 x 10-24 g (approx. the same mass as a hydrogen atom)

• Neutrons have: – no charge

– a mass equal to 1.6749 x 10-24 g

– Neutrons were proposed by Ernest Rutherford in 1907 (to account for a mass discrepancy in the nucleus) and discovered in 1932 by James Chadwick.

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Subatomic Particles -Continued

• Electrons have:

– a charge equal to -1.6022 x 10-19 C

(expressed as -1)

– a mass equal to 9.1094 x 10-28 g (1836 times less than the mass of one

hydrogen atom)

– Electrons were discovered in 1897 by J.J. Thomson.

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Subatomic Particles

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Atomic Number and Mass Number

• Atomic Number (Z)– the number of protons in the nucleus of an element’s

atom– is generally found on the periodic table above the

elemental symbol• Mass Number (A)

– the number of protons and neutrons in the nucleus of an element’s atom

– is generally found below the elemental symbol on the periodic table

A = Z + N• Neutron Number (N)

– the number of neutrons in the nucleus of an element’s atom

N = A - Z2 -

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Atomic Number and Mass Number

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Atomic Mass

Au79

197

110

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Isotopes• An isotope of an element

– is an atom that contains a specific number of neutrons.

– Many elements have multiple isotopes.

– Specific isotopes have many applications, particularly in medical testing, imaging, and treatment.

• An isotope symbol (Nuclide Symbol)

– is a common notation that represents the mass number, atomic number, and elemental symbol.

• The subscript in the isotope symbol is the atomic number.

• The superscript in the isotope symbol is the mass number.

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Isotopes

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Practice – Isotope Symbols

• Practice writing the isotope symbols

for the following isotope pairs.

1. Carbon-13 and carbon-14

2. Chlorine-35 and chlorine-37

3. Uranium-235 and uranium-238

4. Lithium-6 and lithium-7

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Practice Solutions – Isotope

Symbols• Practice writing the isotope symbols for the

following isotope pairs.

1. Carbon-13 and carbon-14

2. Chlorine-35 and chlorine-37

C and C14

6

13

6

Cl and Cl37

17

35

17

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Practice Solutions – Isotope

Symbols• Practice writing the isotope symbols for the

following isotope pairs.

3. Uranium-235 and uranium-238

4. Lithium-6 and lithium-7

U and U238

92

235

92

Li and Li7

3

6

3

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Ions• An ion

– is a charged atom that contains more or less electrons than protons.

– The overall charge is represented as a superscript to the right of the elemental symbol.

• Ions can be classified as cations or anions.

– Cations

• are ions with a positive charge

• have less electrons than protons

– Anions

• are ions with a negative charge

• have more electrons than protons

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Ions

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Practice – Counting Protons and Electrons

• Write the number of protons and

electrons for the following ions:

1. Na+

2. Cl-

3. O2-

4. Al3+

5. P3-

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Practice Solutions – Counting Protons and Electrons

• Write the number of protons and electrons for the following ions:

1. Na+

Sodium has an atomic number equal to 11. Thus, it has 11 protons. It also has a +1 charge, and therefore has 1 less electron than proton. Thus, it has 10 electrons.

2. Cl-

Chlorine has an atomic number equal to 17. It has 17 protons. It also has a -1 charge, and therefore has 1 more electron than proton. Thus, it has 18 electrons.

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Practice Solutions – Counting Protons and Electrons

• Write the number of protons and electrons for the following ions:

3. O2-

Oxygen has an atomic number equal to 8. It has 8 protons. It also has a -2 charge, and therefore has 2 more electrons than protons. Thus, it has 10 electrons.

4. Al3+

Aluminum has an atomic number equal to 13. It has 13 protons. It also has a +3 charge, and therefore has 3 less electrons than protons. Thus, it has 10 electrons.

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Practice Solutions – Counting Protons and Electrons

• Write the number of protons and

electrons for the following ions:

5. P3-

Phosphorus has an atomic number

equal to 15. It has 15 protons. It also has a -3 charge, and therefore has 3

more electrons than protons. Thus, it has 18 electrons.

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Atomic Mass

• Individual atomic masses are determined by mass spectrometry.

• Instead of expressing atomic masses in

grams (a very small number), chemists express atomic masses in atomic mass

units.

– An atomic mass unit (amu) is equal to 1/12 the mass of a carbon-12 atom.

1 amu = 1/12 x mass of 1 12C atom

1 amu = 1.6606 x 10-24 g2 -

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Atomic Mass

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Relative Atomic Mass• Because most elements have multiple

isomers, the masses on the periodic table cannot describe only 1 isotope’s individual atomic mass.

• Therefore, the mass numbers on the periodic table are relative atomic masses:

– Relative atomic mass is the average mass of the individual isotopes of an element, taking into account the naturally occurring relative abundance of each.

– To find the relative atomic mass for an element, sum the mass contributions from each isotope of the element.

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Relative Atomic MassMass contribution from isotope = Isotope mass x relative

abundanceRelative Atomic Mass = Mass contribution from 1st

isotope + Mass contribution from 2nd isotope + …

• Example69.An unknown element (X) discovered on a planet in another

galaxy was found to exist as two isotopes. Their atomic masses and percent abundances are listed in the following table. What is the relative atomic mass of the element?

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25.00

75.00

Natural Abundance (%)

19.996

21.995

Mass (amu)

20X

22X

Isotope

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Mendeleev’s Table• Russian chemist Dmitri Mendeleev

developed and published the basic arrangement of the periodic table between 1869 and 1871.

• Mendeleev arranged the elements in order of increasing relative atomic mass (protons had not been discovered yet). The elements on the modern periodic table are arranged in order of increasing atomic number.

• He also grouped elements with similar properties into columns and rows so that the properties of the elements varied in a regular pattern (periodically).

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Mendeleev’s Table

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The Modern Periodic Table• The elements in the modern periodic table are

arranged by increasing atomic number (Z) and in columns and rows to emphasize periodic properties.

• The columns are collectively called families or groups and are designated in two ways:

1. A Roman numeral (I through VIII) and a letter (A or B)

2. An Arabic number (1-18)

• The rows are collectively called periods and are designated by an Arabic number (1-7).

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The Modern Periodic Table

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Metals, Nonmetals, and Metalloids

• The periodic table has many classifications. Groups and Periods are one classification. Another classification denotes metals,

nonmetals, and metalloids.

– A stair-step line starting at boron (B)

separates metals (to the left of the line) from nonmetals (to the right of the line).

– The metalloids exist along the line.

• Metalloids are elements that have physical

properties resembling a metal, but the chemical reactivity of a nonmetal.

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Metals, Nonmetals, and Metalloids

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Main-Group Elements and Transition Metals

• Main-group elements (also called representative elements) contain any element in the eight groups designated with the letter A. (In the Arabic numbering, groups 1, 2, and 13-18)

• Transition metals contain any element in the 10 groups designated with the letter B. (In the Arabic numbering, groups 3-12)

• Inner-transition metals contain the lanthanides and actinides listed separately at the bottom of the table. (14 Groups)

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Main-Group Elements and Transition Metals

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14

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Common Group Names• Some groups have descriptive names that are

commonly used instead of their group numbers.

– Alkali metals

• Group 1 (IA) metals (hydrogen is a nonmetal)

• are considered reactive because the react

readily with other elements and compounds

– Alkaline earth metals

• Group 2 (IIA) metals

• are more reactive than the transition metals but less reactive than alkali metals

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Common Group Names-Continued

• Some groups have descriptive names that are commonly used instead of their group numbers.

– Halogens (Halides)

• Group 17 (VIIA) nonmetals

• exist naturally as diatomic molecules

– Noble gases

• Group 18 (VIIIA) nonmetals

• are also called inert gases

• are so named because they do not chemically react with other elements (with the exception of krypton and xenon)

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Common Group Names

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Ions and the Periodic Table• The noble gases are the most stable (least

reactive) elements on the periodic table.

– Their stability is associated with the number of electrons they contain (8 electrons in their

outermost layer (or shell)).

– Many atoms in the main-group elements gain

or lose electrons to achieve similar stability.

• Metals tend to lose electrons, and

therefore become cations.

• Nonmetals tend to gain electrons, thereby

becoming anions.

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Ions and the Periodic Table

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Practice – Predicting Charges for Ions

• Write the symbol for the ion that each of the following elements is predicted to

form:

1. Beryllium

2. Aluminum

3. Phosphorus

4. Chlorine

5. Oxygen

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Appendix B in Lab Book - Ions to Learn

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Practice Solutions – Predicting Charges for Ions

• Write the symbol for the ion that each of the following elements is predicted to form:1. Beryllium

Beryllium is in group IIA (2), so it will lose two electrons to form Be2+.

2. Aluminum

Aluminum is in group IIIA (13), so it will lose three electrons to form Al3+.

3. Phosphorus

Phosphorus is in group VA (15), so it will gain three electrons to form P3-.

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Practice Solutions – Predicting Charges for Ions

• Write the symbol for the ion that each of the following elements is predicted to form:

4. Chlorine

Chlorine is in group VIIA (17), so it will gain one electron to form Cl-.

5. Oxygen

Oxygen is in group VIA (16), so it will gain two electrons to form O2-.

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