Elements, Atoms & Ions

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Introductory Chemistry:

A Foundation

FOURTH EDITION

by Steven S. Zumdahl

University of Illinois

Elements, Atoms & Ions

Chapter 4 Elements

• Aims: To learn about the relative abundances of the elements, learn the names of elements

• Over 114 known, of which 88 are found in nature – others are man-made

• Abundance is the percentage found in nature – oxygen most abundant element (by mass) on earth and in

the human body

– the abundance and form of an element varies in different parts of the environment

• Each element has a unique symbol

• The symbol of an element may be one letter or two – if two letters, the second is lower case

Table 4.1: Distribution (Mass Percent) of the 18 Most Abundant Elements in the Earth's Crust,

Oceans, and Atmosphere

Table 4.2: Abundance of elements in the human body

The Symbols for the Elements

• Aim: to learn the names/symbols for some of the elements

• You need to know the elements in Table 4.3

for Quiz #3!!

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Table 4.3: The names and symbols of the most common elements

Dalton’s Atomic Theory

• Aims:

– Learn Dalton’s Atomic Theory – Learn Law of Constant Composition

– In the 18^thcentury, scientists studying the nature of things agreed upon the following things:

• Most natural materials are mixtures of pure substances

• Pure substances are either elements or combinations of elements

• A given compound always contains the same proportions (by mass) of the elements. e.g., water always contains 8 g of oxygen for every 1 g of hydrogen. This principle is known as the Law of Constant Composition. It means a given compound always has the same composition.

Figure 4.1: John Dalton (1766-1844)

An English scientist and teacher was aware of these observations and formulated an explanation.

His explanation became known as Dalton’s Atomic Theory

Dalton’s Atomic Theory

1. Elements are composed of atoms – tiny, hard, unbreakable, spheres

2. All atoms of a given element are identical – all carbon atoms have the same chemical and physical

properties

3. Atoms of a given element are different from those of any other element

– carbon atoms have different chemical and physical properties than sulfur atoms

Dalton’s Atomic Theory

4. Atoms of one element combine with atoms of other elements to form compounds.

– Law of Constant Composition

• all samples of a compound contain the same proportions (by mass) of the elements

– Chemical Formulas

• Describe the proportions of elements in a compound

Dalton’s Atomic Theory

5. Atoms are indivisible in a chemical process.

– all atoms present at beginning are present at the end – atoms are not created or destroyed, just rearranged – atoms of one element cannot change into atoms of

another element

• cannot turn Lead into Gold by a chemical reaction

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Formulas

• Aims:

– Learn how a formula describes a compounds composition – Learn how to write formulas

• a compound is a distinct substance that is composed of atoms of two or more elements

• Formulas describe the compound by describing the number and type of each atom in the simplest unit of the compound

– molecules or ions

Formulas Describe Compounds

Rules for Writing Formulas

• each element represented by its letter symbol

• the number of atoms of each element is written to the right of the element as a subscript

– if there is only one atom, the 1 subscript is not written

• polyatomic groups are placed in parentheses – if more than one

• If subscript is one (1), then it is not written

Figure 4.2: Dalton pictured compounds as collections of atmosphere NO, NO₂, and N₂O are represented

Structure of the Atom

• Aims:

– Learn the internal parts of the atom – Understand Rutherford’s experiment to

characterize the atom’s structure

• Dalton’s theory explained compounds as a collection of atoms. But what were the atoms themselves like?

The Nature of the Atom

• Many scientists pondered the nature of the atom in the 1800s

• Physicist J.J. Thomson showed that atoms of any kind can emit tiny negative particles. Therefore all atoms must contain these tiny, negative particle known as electrons

• Since J.J. Thomson knew that atoms are neither positively or negatively charged, he concluded that there must also be positive particles present in the atom to balance the charge to 0 (zero)

Are Atoms Really Unbreakable?

• J.J. Thomson investigated a beam called a cathode ray

• he determined that the ray was made of tiny negatively charged particles we call electrons

• his measurements led him to conclude that these electrons were smaller than a hydrogen atom

• if electrons are smaller than atoms, they must be pieces of atoms

• if atoms have pieces, they must be breakable

• Thomson also found that atoms of different elements all produced these same electrons

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The Electron

• Tiny, negatively charged particle

• Very light compared to mass of atom

– 1/1836^ththe mass of a H atom

• Move very rapidly within the atom

Thomson’s Plum Pudding Model

Thomson concluded from his studies:

1. Atom breakable!!

2. Atom has structure

3. Electrons suspended in a positively charged electric field – must have positive charge to balance negative charge

of electrons and make the atom neutral 4. mass of atom due to electrons

5. atom mostly “empty” space

– compared size of electron to size of atom

Figure 4.3: The plum pudding

model

Figure 4.4: Ernest Rutherford (1871-1937)

A physicist, who found that something was deflecting the α-particles he was studying

He set up an experiment to find out what it was…

Figure 4.5: Rutherford’s experiment on α- particle bombardment of metal foil

•α-particles have a mass of 4 amu & charge of +2 c.u.

• gold has a mass of 197 amu & is very malleable

Rutherford’s Gold Foil Expt

• How can you prove something is empty?

– The Plum Pudding model postulated that the atom was an empty cloud of positive charge with electrons scattered through it

• If the Plum Pudding Model was correct, the α-particles would fly right through his gold foil (like a bullet through a piece of paper)

– Rutherford expected the α-particles to fly through the foil with at most a minor deflection

• But some particles experienced large deflections!!

• Therefore the Plum Pudding Model is not correct!

– The large deflections were due to positive particles hitting a positively charged nucleus

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Rutherford’s Results

• Over 98% of the α particles went straight through

• About 2% of the α particles went through but were deflected by large angles

• About 0.01% of the α particles bounced off the gold foil

Figure 4.6: (a) The results that the metal foil experiment would have yielded if the plum pudding

model had been correct; (b) Actual results

Rutherford’s Nuclear Model

1. The atom contains a tiny dense center called the nucleus

– the volume is about 1/10 trillionth the volume of the atom

2. The nucleus is essentially the entire mass of the atom

3. The nucleus is positively charged

– the amount of positive charge of the nucleus balances the negative charge of the electrons 4. The electrons move around in the empty space of

the atom surrounding the nucleus

Structure of the Nucleus

• The nucleus was found to be composed of two kinds of particles

• Some of these particles are called protons – charge = +1

– mass is about the same as a hydrogen atom

• Since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons

• The other particle is called a neutron – has no charge

– has a mass slightly more than a proton

The Modern Concept of the Atom

• Aim:

– To understand the main features of subatomic particles

0 (none) 1839

Neutron

1+

1836 Proton

1- 1

Electron

Relative Charge Relative Mass

Particle

The Modern Atom

• We know atoms are composed of three main pieces - protons, neutrons and electrons

• The nucleus contains protons and neutrons

• The nucleus is only about 10

^-13

cm in diameter

• The electrons move outside the nucleus with an average distance of about 10

^-8

cm

– therefore the radius of the atom is about 10⁵times larger than the radius of the nucleus

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Figure 4.9: A nuclear atom viewed in cross

section

Components of Atoms

• If all atoms are composed of the same components (electrons, protons, and neutrons) …Why do different atoms have different chemical properties??

• The answer is the number and arrangement of the electrons

– Electrons account for most of the “volume” of an atom – Electrons are the part of the atom that intermingle with

other atoms, so the number and arrangement of electrons affect chemical behavior

Bohr’s Model

• We’ll Talk about this more in Chapter 10…

• Planetary Model

– Based upon the orbits of our solar system.

• Not a 100% correct model, but good enough to explain some concepts.

The Orbits or Shells (Bohr)

• Also talk more about this in Chapter 10

• So there are these orbits around the nucleus of the atom where the electrons are.

• The first orbit can only hold 2 electrons.

• All other orbits can hold 8 electrons.

– Octet Rule

• Atoms will arrange themselves in order to obtain 8 electrons.

• Except: hydrogen and helium

• Valence Orbit: the outer most orbit

– This is the orbit that is used to create ion and is used in bonding.

– The electrons in this orbit are called valence shell electrons.

– So it is kind of important…

• Bohr’s model can only be used for the first 3 periods in the periodic table.

Isotopes

• Aims:

• Learn the following terms

– Isotope

– Atomic number – Mass number

– Understand the symbols used to describe atoms

Isotopes

• All atoms of an element have the same number of protons

• The number of protons in an atom of a given element is the same as the atomic number

– found on the Periodic Table

• Atoms of an element with different numbers of neutrons are called isotopes

• All isotopes of an element are chemically identical – undergo the exact same chemical reactions

• Isotopes of an element have different masses

• Isotopes are identified by their mass numbers – mass number = protons + neutrons

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Figure 4.10: Two isotopes of sodium

Mass Number = # protons + # Neutrons Atomic number = # protons

Symbols and Calculations

• You should be able to do the following:

• Interpret a symbol for an isotope (Ex 4.2)

• Write a symbol for an isotope (Ex 4.3)

• Calculate a mass number (Ex 4.4)

Elements

• Arranged in a pattern called the Periodic Table

• Position on the table allows us to predict properties of the element

• Metals

– about 75% of all the elements

– lustrous, malleable, ductile, conduct heat and electricity

• Nonmetals

– dull, brittle, insulators

• Metalloids

– also know as semi-metals

– some properties of both metals & nonmetals

Figure 4.11: The periodic table

The Modern Periodic Table

• Elements with similar chemical and physical properties are in the same column

• Columns are called Groups or Families

• Rows are called Periods

• Each period shows the pattern of properties repeated in the next period

Figure 4.11: The periodic table

Group

Period

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The Modern Periodic Table

• Main Group = Representative Elements

– “A” columns

• Transition Elements

– all metals

• Bottom rows = Inner Transition Elements = Rare Earth Elements

– metals

– really belong in Period 6 & 7

Figure 4.12: The elements classified as metals and nonmetals

Metalloids

Important Groups

• Group 8 = Noble Gases

• He, Ne, Ar, Kr, Xe, Rn

• all colorless gases at room temperature

• very non-reactive, practically inert

• found in nature as a collection of separate atoms uncombined with other atoms

• Noble Metals

• Ag, Au, Pt

• all solids at room temperature

• least reactive metals

• found in nature uncombined with other atoms

Figure 4.13:

Argon gas consists of a collection of separate argon

atoms

Figure 4.14: Gaseous nitrogen and oxygen contain diatomic (two-atom) molecules

Figure 4.15: The decomposition of two water molecules (H₂O) to form two hydrogen molecules (H₂) and an oxygen molecule (O₂)

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Important Groups - Halogens

• Group 7A = Halogens

• very reactive nonmetals

• react with metals to form ionic compounds

• HX all acids

• Fluorine = F₂ – pale yellow gas

• Chlorine = Cl₂ – pale green gas

• Bromine = Br₂

– brown liquid that has lots of brown vapor over it – Only other liquid element at

room conditions is the metal Hg

• Iodine = I₂ – lustrous, purple solid

Figure 4.16: (a) Sodium chloride (table salt) can be decomposed to the elements sodium metal and

chlorine gas (b)

Allotropes

• Many solid nonmetallic elements can exist in different forms with different physical properties, these are called allotropes

• the different physical properties arise from the different arrangements of the atoms in the solid

• Allotropes of Carbon include

– diamond

– graphite

– buckminsterfullerene

Figure 4.17: In solid metals, the spherical atoms are packed closely together

Figure 4.18a: The three solid elemental (allotropes) forms of carbon

Figure 4.18b: The three solid elemental (allotropes) forms of

carbon

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Figure 4.18c: The three solid elemental (allotropes) forms of carbon

(c) Buckminsterfullerene

Electrical Nature of Matter

• Most common pure substances are very poor conductors of electricity

– with the exception of metals and graphite – Water is a very poor electrical conductor

• Some substances dissolve in water to form a solution that conducts well - these are called electrolytes

• When dissolved in water, electrolyte compounds break up into component ions

– ions are atoms or groups of atoms that have an electrical charge

Figure 4.20: (a) Pure water does not conduct a current; (b) Water containing a dissolved salt

conducts electricity

Ions

• ions that have a positive charge are called cations

– form when an atom loses electrons

• ions that have a negative charge are called anions

– form when an atom gains electrons

• ions with opposite charges attract

– therefore cations and anions attract each other

• moving ions conduct electricity

• compound must have no total charge, therefore we must balance the numbers of cations and anions in a compound to get 0 total charge

Figure 4.19: The ions formed by selected members of Groups 1, 2, 3, 6, and 7

Atomic Structures of Ions

• Metals form cations

• For each positive charge the ion has 1 less electron than the neutral atom

– Na = 11 e^-, Na⁺= 10 e^- – Ca = 20 e^-, Ca⁺²= 18 e^-

• Cations are named the same as the metal

sodium Na → Na⁺+ 1e^- sodium ion calcium Ca → Ca⁺²+ 2e^- calcium ion

• The charge on a cation can be determined from the Group number on the Periodic Table for Groups IA, IIA, IIIA

– Group 1A ⇒ +1, Group 2A ⇒ +2, (Al, Ga, In) ⇒ +3

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Figure 4.21a: The arrangement of sodium ions (Na⁺)

and chloride ions (Cl^-) in the ionic compound sodium

chloride.

Figure 4.21b: Solid sodium chloride highly magnified.

Atomic Structures of Ions

• Nonmetals form anions

• For each negative charge the ion has 1 more electron than the neutral atom

– F = 9 e^-, F^-= 10 e^- – P = 15 e^-, P^3-= 18 e^-

• Anions are named by changing the ending of the name to -ide

fluorine F + 1e^-→ F^-fluoride ion oxygen O + 2e^-→ O^2- oxide ion

• The charge on an anion can be determined from the Group number on the Periodic Table

– Group 7A ⇒ -1, Group 6A ⇒ -2