Chapter 4 Elements, Atoms, and Ions. Ch#4 Contents 4.1The Elements 4.2Symbols for the Elements...

Chapter 4Elements, Atoms, and Ions

Ch#4 Contents4.1 The Elements4.2 Symbols for the Elements4.3 Dalton’s Atomic Theory4.4 Formulas of Compounds4.5 The Structure of the Atom4.6 Modern Concept of Atomic Structure4.7 Isotopes4.8 Introduction to the Periodic Table4.9 Natural States of the Elements4.10 Ions4.11 Compounds that Contain Ions

ElementsA substance that cannot be broken down by chemical means. Elements are defined by the number of protons they possess. •115 known: 88 found in nature, others are man made.•Most of the know elements are found on the Periodic Chart handing on the wall.•One or two letters are used to represent an element and they are called the symbol•Some elements are solids (black), some are liquids (blue) and some are gases (red)

Element Distribution in Earth’s Surface

Symbols• Each element has a unique one- or two-letter

symbol.• First letter is always capitalized and the second is

not.• The symbol usually consists of the first one or two

letters of the element’s name. Examples:

Oxygen O Krypton Kr

• Sometimes the symbol is taken from the element’s original Latin, Greek, English, or German name.

Examples: Gold Au aurum Lead Pb plumbum Wolfram W tungsten

Symbols

Located on the front page of your textbook

Dalton’s Atomic TheoryIn 1808 Dalton published A New System of Chemical Philosophy, where he presented his theory of atoms:•Each element is made up of tiny particles called atoms.•The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways•Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.•Chemical changes involve reorganization of the atoms, to different ratios.

Dalton’s Atomic TheoryWhich of the following statements regarding Dalton’s atomic theory are still believed to be true?•Elements are made of tiny particles called atoms. • All atoms of a given element are identical. • A given compound always has the same relative numbers and types of atoms. • Atoms are indestructible.

Dalton’s Atomic TheoryWhich of the following statements regarding Dalton’s

atomic theory are still believed to be true?• Elements are made of tiny particles called atoms. • All atoms of a given element are identical. • A given compound always has the same relative

numbers and types of atoms. • Atoms are indestructible.

Mendeleev’s Periodic Table• Dmitrii I. Mendeleev

arranged elements in the periodic table by their chemical and physical properties.

• He left open spaces in his periodic table to account for elements not yet discovered.

The Modern Periodic Table• The modern periodic table is also based on a

classification of elements in terms of their physical and chemical properties.

• The horizontal rows are called periods.• Columns contain elements of the same family or

group.• Transition metals are the elements in group 3

through 12 in the periodic table.

Formulas• Compound – distinct substance that is composed of

the atoms of two or more elements and always contains exactly the same whole number ratio of those elements.

• Chemical Formulas – expresses the types of atoms and the number of each type in each formula unit, or molecule of a given compound.

• Formula Examples– H2O (molecule)

– C6H12O6 (molecule)

– NaCl (Formula Unit)

Metals

Atomic Structure

• Postulated the existence of electrons using cathode-ray tubes.

• The atom must also contain positive particles that balance exactly the negative charge carried by particles that we now call electrons.

J. J. Thomson (1898—1903)

Cathode-ray Tube

Atomic Structure

Reasoned that the atom might be thought of as a uniform “pudding” of positive charge with enough negative electrons scattered within to counterbalance that positive charge.

William Thomson (Plum Pudding Model)

Atomic Structure

• Explained the nuclear atom.• Atom has a dense center of positive charge called

the nucleus.• Electrons travel around the nucleus at a relatively

large distance.• A proton has the same magnitude of charge as the

electron, but its charge is positive.

Ernest Rutherford (1911)

Atomic Structure

• Most nuclei also contain a neutral particle called the neutron.

• A neutron is slightly more massive than a proton but has no charge.

Rutherford and Chadwick (1932)

Atomic Structure• Electrons – found

outside the nucleus; negatively charged

• Protons – found in the nucleus; positive charge equal in magnitude to the electron’s negative charge

• Neutrons – found in the nucleus; no charge; virtually same mass as a proton

Atomic Structure• The nucleus is:

– Small compared with the overall size of the atom.– Extremely dense; accounts for almost all of the

atom’s mass.

Atomic Structure

Atomic StructureWhy do different atoms have different chemical properties?• The chemistry of an atom arises from its

electrons.

• Electrons are the parts of atoms that “intermingle” when atoms combine to form molecules.

• It is the number of valence (furthest from nucleus) electrons that really determines chemical behavior.

Isotopes• Atoms with the same number of protons but

different numbers of neutrons.• Show almost identical chemical properties;

chemistry of atom is due to its electrons.• In nature most elements contain mixtures of

isotopes.

Symbols of Isotopes• Each element consists of atoms with the same

number of protons in the nucleus. This number is called atomic number (Z).

• Protons and neutrons in atomic nuclei are called nucleons.

• The mass number (A) is the total number nucleons in one atom of an element.

Writing the Symbol of an Isotope

• A is the mass number

• Z is the atomic number

• X is the atomic symbol

Isotopes

• C = the symbol for carbon

• 6 = the atomic number (6 protons)

• 14 = the mass number (6 protons and 8 neutrons)

• C = the symbol for carbon• 6 = the atomic number (6

protons)• 12 = the mass number (6

protons and 6 neutrons)

Two Isotopes of Sodium

BOHR MODELS OF HYDROGEN(The three isotopes of hydrogen)

Hydrogen-1 Hydrogen-2 Hydrogen-3

BOHR MODELS OF HYDROGEN

H HH1 2 3

1 1 1Atomic Number (number of protons); found on the periodic chart

H HH1 2 3

1 1 1Atomic Number (number of protons)

Mass number (sum of protons and neutrons); Not found on periodic chart

H HH1 2 3

1 1 1Atomic Number (number of protons)

Mass number (sum of protons and neutrons); Not found on periodic chart

Oxidation number (Protons – electrons)

Average Atomic Masses• A mass spectrometer is an instrument that

measures precise masses and relative amounts of ions of atoms and molecule.

• The natural abundance of an isotope is its relative proportion among all the isotopes found a natural sample.

• The average atomic mass of an element is calculated by multiplying the natural abundance of each isotope by its exact mass in atomic mass units and then summing these products.

Relative Atomic Mass of HydrogenWhen computing the average atomic mass of any element the radioactive (unstable) isotopes are excluded since there relative abundances are slowly decreasing.

Hydrogen has three isotopes. The first two protium and deuterium are stable isotopes and the third tritium is unstable, thus excluded in the calculation.

The relative abundance of isotopes of a particular element is constant here on our planet Earth. Mass spectroscopy gives information of relative abundances and relative masses of isotopes.

Some atoms naturally occur as a mixture of isotopes.

e.g. 11H, 1

2H and 13H

The atomic weights given in the periodic table take this into account.

They are the average atomic weight taking into account the amounts of each isotope present.

Relative Atomic Mass

Mass spectroscopy gives the following information for the two stable isotopes of hydrogen.

isotope name relative abundance Mass (amu) Hydrogen-1(protium) 99.985 % 1.007825

Hydrogen-2(deuterium) 0.015 % 2.01355

To calculate a weighted average, convert the percent to a decimal by moving the decimal two places to the left. Then multiply this number by the mass and add the two masses.

Relative Atomic Mass of Hydrogen

0.99985 X 1.007825 = 1.007673

0.0015 X 2.01355 = 0.003021

Relative Atomic Mass of Hydrogen

1.010694 amu

1.0107 amu

Example

Isotope Mass (amu) Natural abundance (%)

Neon-20 19.9924 90.4838

Neon-21 20.99395 0.2696

Neon-22 21.9914 9.2465

Neon is found in three isotopes in nature.

19.9924 x 0.904838 = 20.99395 x 0.002696 =21.9914 x 0.092465 =

18.0898323 amu0.056599689 amu2.033434801 amu

20.1797 amu

If we assign a mass of 12 atomic mass units (u or sometimes amu) to a carbon-12 atom then we can compute the relative atomic weight for any other atom.

Atomic weights or atomic masses are given in your periodic table.

Relative Atomic Mass

Natural State of Elements• Most elements are very reactive.• Elements are not generally found in uncombined

form. Exceptions are:

• Noble metals – gold, platinum and silver• Noble gases – Group 8

Diatomic Molecules

Allotropes• Different forms of a given element.• Example:

Solid carbon occurs in three forms.• Diamond• Graphite• Buckminsterfullerene

Carbon Allotropes

Ions• Atoms can form ions by gaining or losing electrons.

Metals tend to lose one or more electrons to form positive ions called cations and are named by using the name of the parent atom.

Nonmetals tend to gain electrons to form negative ions called anions and are named by using the root of the atom name followed by the suffix –ide.

• Atoms become more stable by gaining or losing electrons.

Cations are generally named by using the name of the parent atom.

Common Oxidation States

By Group Number

Sample ProblemAn ion with a 3+ charge contains 23 electrons. Which ion is it?

Sample ProblemAn ion with a 3+ charge contains 23 electrons. Which ion is it? Fe3+

A certain ion X+ contains 54 electrons and 78 neutrons. What is the mass number of this ion?

A certain ion X+ contains 54 electrons and 78 neutrons. What is the mass number of this ion? 133

Ionic Compounds• Ions combine to form ionic compounds.• The simplest ratio between the ions is called the

empirical formula or a formula unit. • Properties of ionic compounds

– High melting points– Conduct electricity

• If melted• If dissolved in water

• Ionic compounds are electrically neutral.• The charges on the anions and cations in the

compound must sum to zero.

Ionic compounds do not exist as discrete molecules. Instead they exist as crystals where ions of opposite charges occupy positions known as lattice sites.

Ions combine in the ratio that results in zero charge to form ionic compounds. Which ions are the smaller ones?

Crystal Lattice of NaCl

Ionic Compounds

Formula Writing RulesStep 1 Write the symbols of the elements.

Step 2 Assign oxidation numbers (O.N. = P – e)

Step 3 Slide with Clyde! (number only)

Step 4 Reduce if the compound is ionic

ExamplesWrite formulas for the following two elements

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Calcium and phosphorus

Step 2 Assign oxidation numbers O.N. = (P – e)

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Na O2≡

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Ca O2 2 CaOreduced

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Ca O2 2 CaOreduced

Sodium and oxygen

Sodium and nitrogen

Calcium and oxygen

Ca O2 2 CaOreduced

Ca P3 2

The End

Chapter 4 Elements, Atoms, and Ions. Ch#4 Contents 4.1The Elements 4.2Symbols for the Elements...

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