Elements, Atoms, and Ions€¦ · 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types

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Sec. 0.3: Chemical Foundations

Elements, Atoms, and Ions

Objectives • Learn the names and symbols for some elements.

• Learn about the relative abundance for some elements.

• Learn about Dalton’s theory of atoms.

• Understand the law of constant composition.

• Learn about how a formula describes a compound’s composition.

• Understand Rutherford’s experiment and its impact on atomic structure.

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Objectives • Describe important features of subatomic particles.

• Learn about isotope, atomic number and mass number.

• Understand the use of the symbol X.

• Learn the various features of the periodic table.

• Learn the properties of metals, nonmetals, and metalloids.

• Describe the formation of ions from their parent atoms.

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Objectives • Predict which ion a given element forms by using the periodic table.

• Describe how ions combine to form neutral compounds.

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Section 3.1: The Elements

• Remember, elements are combined to form molecules the way letters are combined to form words.

• Presently there are about 115 known elements.

• Only 88 occur naturally, the rest are made in laboratories.

• Only 9 elements account for most of the compounds found in the Earth’s crust.

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The Elements • Scientists use the word element in many different ways.

• Sometimes it is referred to in the microscopic sense: – A single atom of Au or Ag could be referred to as an element.

– Also, molecules such as O2 or N2, are referred to as elements.

• In the macroscopic sense we can refer to a bar of “pure” iron or a 24k gold ring as elements.

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The Elements • When we say something contains a particular element we do not necessarily mean free atoms, but may also mean in a form combined with other elements in some compound.

• Our bodies contain many “trace” elements – elements that are present in very small amounts, but are crucial to life.

• Some of these elements include: arsenic, chromium, cobalt, copper, fluorine, iodine, manganese, molybdenum, nickel, selenium, silicon and vanadium. 3-7

Section 3.2: Symbols For The Elements

• Just as each state has a two-letter abbreviation, each element has a one- or two-letter symbol to make life simple for chemists.

• The list of trace elements from the previous slide can be simplified to: As, Cr, Co, Cu, F, I, Mn, Mo, Ni, Se, Si, & V.

• Notice the first letter is ALWAYS capitalized and the second letter, if present, is NEVER capitalized.

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Symbols For The Elements • Some symbols make sense like O for oxygen and H for hydrogen or Ni for nickel.

• Others, like Pb for lead or Fe for iron, don’t automatically make sense; they originated from the Greek or Latin names of plumbum (Pb) and ferrum (Fe).

• The only real way to learn them all is to memorize them. Chemists have arranged them in a periodic table to help with that.

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While there are well over 100 different

elements, many are

fairly rare; we should know

the most common

elements.

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Section 3.3: Dalton’s Atomic Theory

• Scientists studying matter in the eighteenth century made the following observations: –Most natural materials are mixtures of pure substances.

– Pure substances are either elements or combinations of elements called compounds.

– A given compound always contains the same proportions (by mass) of the elements.

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Dalton’s Atomic Theory • One particular English scientist named John Dalton attempted to explain these observations in 1808.

• Dalton made his living as a teacher in Manchester, England; he started a school in his town when he was just 12 years old!

• He never married, was colorblind (to red) and liked to bowl every Thursday afternoon.

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Dalton’s Atomic Theory

1. Elements are made of tiny particles called atoms.

2. All atoms of a given element are identical.

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

4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms.

5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way atoms are grouped together. 3-13

Section 3.4: Formulas of Compounds

• The types of atoms and the number of each type in each unit (molecule) of a given compound are conveniently expressed by a chemical formula.

• The atoms are indicated by their symbols and the number of each type is indicated by a subscript (unless there is only one). – Ex) C6H12O6 or H3PO4

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Practice • Write the formula for each of the

following compounds, listing the elements in the order given: a. A molecule contains four phosphorous atoms

and ten oxygen atoms.

b. A molecule contains one uranium atom and six fluorine atoms.

c. A molecule contains one aluminum atom and three chlorine atoms.

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Section 3.5: The Structure of the Atom

• In the late 1890’s J.J. Thomson, an English physicist, determined atoms of any element could be made to emit tiny negative particles.

• He showed they were repelled by the negative part of an electric field.

• He had discovered electrons. • He also concluded atoms must contain positive charges to cancel out the negative (atoms are electrically neutral).

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The Structure of the Atom • Building on J.J. Thomson’s discoveries, Lord Kelvin proposed a “plum pudding” model.

• Consider pudding with raisins in it. Now, imagine the raisins are negatively-charged and the pudding is positively-charged.

• The positive charge was cancelled out by the negative charges giving an overall charge of zero.

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Figure 3.3: Plum Pudding model of an atom.

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The Structure of the Atom • In 1911 another physicist named Ernest Rutherford performed his famous “Gold Foil Experiment” that concluded the plum pudding model could not be correct.

• His experiment involved firing alpha particles at a sheet of thin gold foil.

• The foil was surrounded by a detector coated with a substance that produced tiny flashes when hit by an alpha particle.

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Figure 3.5: Rutherford’s experiment.

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The Structure of the Atom • Since alpha particles are about 7500 times more massive than electrons, Rutherford expected them to tear through the foil the way a bullet would go through paper.

• What actually happened was most passed straight through, but some were deflected at great angles and even reflected backwards!

• According to the plum pudding model he expected ALL to pass through with a few being deflected VERY slightly. 3-21

Figure 3.6: Results of foil experiment if

Plum Pudding model had been correct.

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The Structure of the Atom • He concluded the positively-charged alpha particles could have only been deflected by a dense center of concentrated positive charge.

• He came up with the model of a nuclear atom: a nucleus composed of positively-charged protons orbited by negatively-charged electron made up of mostly empty space.

• He eventually came up with idea of neutrons also residing in the nucleus.

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Figure 3.6: Actual results.

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The Structure of the Atom • Ultimately, Rutherford’s model gave us the three subatomic particles: protons (charge = +1) and neutrons (charge = 0) in the small central nucleus and tiny electrons (charge = -1) orbiting the nucleus.

• If an atom were blown up to the size of a professional football stadium, the nucleus would be the size of a fly on the 50-yard line.

• If the nucleus were the size of a grape, the atom’s radius would be about a mile. 3-25

The Structure of the Atom • J.J. Thomson discovered electrons using a cathode ray tube (CRT).

• CRT’s are sealed glass tubes containing a gas and a metal plate at each end connected to external wires.

• When an electric current was applied to the plates a glowing beam was produced; he was convinced it was a stream of electrons.

• This technology is still used today in televisions and computer monitors.

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Figure 3.7: Schematic of a cathode ray tube.

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Section 3.6: The Modern Concept of Atomic Structure

• Today, the view of the atom is: – A tiny nucleus about 10-13 cm in diameter.

– Electrons that move around the nucleus at an average distance of about 10-8 cm away.

– Electrons and protons having equal and opposite charges while neutrons have no charge.

– Protons and neutrons almost 2000 times more massive than electrons.

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Figure 3.9: A nuclear atom viewed in cross

section.

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Modern Atomic Structure • Every atom is composed of the three basic subatomic particles.

• Different elements have different numbers of each of these particles.

• The reason one element behaves differently than another lies in the number and arrangement of their electrons.

• When atoms get close to each other their electron “clouds” can overlap and interact.

• We’ll learn more about this later. 3-30

Section 3.7: Isotopes • We now know each element has a unique number of protons and electrons (they must be equal), but what about the number of neutrons?

• Dalton assumed any two atoms of a given element were identical; not quite correct.

• We can have two atoms of the same element (same number of protons) with different numbers of neutrons.

• These are called isotopes. 3-31

Figure 3.10: Two isotopes of sodium.

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Isotopes • There are two important numbers

associated with any given element: 1. Atomic Number – The number of protons in a

nucleus.

2. Mass Number – The SUM of the number of protons AND neutrons (a.k.a. nucleons) in a nucleus (NOT the sum of their masses).

• We should note that two different isotopes will have the same atomic number, but different mass numbers.

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Isotopes • Scientists like to use symbols as shorthand for

these terms: – X = the symbol of the element

– A = the mass number (nucleons)

– Z = the atomic number (protons)

• A generic representation of any given element would look as follows:

Z

A X3-34

Isotopes • The two previous examples of isotopes of sodium would be:

11

23Na 11

24Na

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•The example on the left would contain 11 protons and 12 neutrons (23-11=12). •The example on the right would contain 11 protons and 13 neutrons (24-11=13).

Practice Problems • Write the symbol for each of the

following atoms, and list the number of protons, neutrons, and electrons for each. 1) The cesium atom with a mass number of 132.

2) The iron atom with a mass number of 56.

3) The krypton atom that has 48 neutrons.

4) The nitrogen atom that has 6 neutrons.

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Figure 3.11: The periodic table

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Figure 3.12: Elements classified as metals and nonmetals.

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Figure 3.13: A collection of argon atoms.

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Figure 3.14: Nitrogen gas contains NXN molecules.

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Figure 3.14: Oxygen gas contains OXO molecules.

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Table 3.5

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Figure 3.15: The decomposition of two water molecules.

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Figure 3.17: Spherical atoms packed closely together.

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Figure 3.19: The ions formed by selected members of groups 1, 2, 3, 6, and 7.

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Section 3.8: Introduction to the Periodic Table

Objectives: To learn about various features of the periodic table. To learn some of the properties of metals, nonmetals, and metalloids.

A Simple Version of the Periodic Table

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• In any box on the Periodic Table, what information can you find?

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C 6

12.01

Atomic number = number of

protons, unique for every

element, no 2 elements have

the same atomic #

Element symbol = can be

1,2 or 3 letters, first letter is

always capitalized, and

succeeding letters are always

lower case Average Atomic Mass = the

weighted average of all the

mass numbers for each

isotope of the element

Weighted Average Atomic Mass

• Remember elements can have different isotopes which means that they vary in their number of neutrons.

• If you have 3 different isotopes of the same element:

– 15 atoms have a mass of 21



We can calculate the weighted average by multiplying the number of atoms by their mass:

(15) (21) = 315 537 = 21.48

(8) (23) = 184 25

(2) (19) =+ 38 average atomic mass

537

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• If you have 10,000 atoms of Cl



What is the average atomic mass of Cl?

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• (7577) (35) = 265195

(2423) (37) = 89651

354528

354528 = 35.45

10,000

Using % to find Average Atomic Mass

• Usually we only know the percents of various isotopes that make up different elements, we can use this to calculate the average atomic mass.

• If we have 100% chlorine:

75.77% of mass is 34.969 -> .7577x34.969 = 26.469

24.23% of mass is 36.966 -> .2423x36.966 = 8.95686

Add the 2 together to get the atomic mass:

26.469 + 8.95686 = 35.45 3-51

Practice

• Oxygen has 3 isotopes 16O, 17O, 18O

99.76% of mass is 16O



What is the average atomic mass?

• Find the atomic mass if 99.64% of mass is 14N and 0.36% is 15N.

• Magnesium has three isotopes. 78.99% magnesium 24 with a mass of 23.9850 amu, 10.00% magnesium 25 with a mass of 24.9858 amu, and the rest magnesium 25 with a mass of 25.9826 amu. What is the atomic mass of magnesium?

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Periodic Table

• When looking at periodic table elements are arranged in horizontal rows by increasing atomic number.

• Horizontal rows are called “Periods”

Periods go left to right

As you move across the period the number of valence electrons increases

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Periodic Table

• The vertical columns are called “Groups” or “Families”

• Elements in families share similar properties

• Each shares the same number of valence electrons in outermost shell

• Can determine the number of valence electrons by the number of the group

• Can use the group number and valence electrons to find the charges of many elements

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Metals, Semimetals, Non-Metals

All elements on the periodic table are grouped as metals, semimetals or metalloids, or non-metals. Due to the arrangement of the periodic table, it is easy to identify each type of element.

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Metals: Fall to left and under the stairs

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Properties of Metals:

Efficient conduction of heat and electricity

Malleability

Ductility

A lustrous appearance

Positively charged ions

Non-Metals: Right and above the stairs

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• Dull, Brittle

• Negatively charged ions

• Nonconductors

-insulators

Semimetals or Metalloids: Makeup the stairs

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• Properties of both metals and non-

metals

• Semiconductors

Lanthanide and Actinide Series • Mostly human made elements

• Radioactive elements

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Group 1A – Alkali Metals

• Members of the group 1A

• Have a +1 charge as ions

• Very reactive with water (stored in oil)

• Extremely malleable

and soft

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Group 1A – Alkali Metals • All members react with water to produce same ratio

of products:

Li + HOH -> LiOH + H2(g)

Na + HOH -> NaOH + H2(g)

K + HOH -> KOH + H2(g)

Rb + HOH -> RbOH + H2(g)

Cs + HOH -> CsOH + H2(g)

• As you move down the group:

– Reactivity with water increases

– Melting points and boiling points decrease

– Density increases

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Group 2A – Alkaline Earth Metals

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• Members of Group 2A

• Form +2 charges as ions

• React with water

Group 2A – Alkaline Earth Metals • The members of this group also react to form the

same ratio:

Mg + HOH -> Mg(OH)2 + H2

Ca + HOH -> Ca(OH)2 + H2

Sr + HOH -> Sr(OH)2 + H2

• As you move down the family we see some patterns once again: – Reactivity increases

– Melting/Boiling point decreases

– Density increases

– Slower to react than alkali metals…don’t have to store them in oil

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Trends for Group 1A & 2A

• Melting/Boiling points increase

• Reactivity increases

• Density increases

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+1

+2

Group 3A or 13 • Form +3 charge as ions

• React with water in a 1metal:3 OH ratio

• Slow reaction with water

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Group 3A or 13

• When looking at Periodic Tables you will see this group numbered 2 different ways, they are still the same group that carries a charge of +3.

• Aluminum is a member of this family. It is a metal, it is reactive. We don’t find pure aluminum in nature, it is always coated with aluminum oxide.

• Aluminum foil is coated with thin layer of vegetable oil to keep shiny

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Groups 4A,14 and 5A,15 • These groups have metals, nonmetals, and

metalloids which makes their properties more difficult to group.

• Group 4A can have a +4

or -4 charge

• Group 5A can have a -3 charge

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Group 6A or 16: The Oxygen Family

• Form -2 charge as ions

• Members below oxygen form oxides with putrid odor:

SO2, SeO2, & TeO2

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Group 7A or 17: Halogen Family

• Exist as diatomic molecules

• Form ions with a -1 charge

• Non-metals

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Group 7A or 17: Halogens or Hydrogen Family

• Sometimes periodic tables will show hydrogen in this family because it’s property match better than the Alkali metals.

• Members of this family are diatomic, they are found as 2 bonded atoms: H2, F2, Cl2, Br2, I2

• Members of this family are toxic and poisonous (exception is H2)

• Ions of this family have -1 charge and are called halides

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• F2 is toxic, and the most reactive element on the periodic table, it can explode almost instantly with anything it mixes with

• Cl2 is a poisonous gas, known as mustard gas in WWII • Br2 is a red liquid with low vapor pressure, so it has a low

boiling point • I2 is a silvery solid that sublimates from solid to gas.

Creates a purple toxic gas • At2 radioactive and very rare, less than 30 grams on Earth,

not studied very often.

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• Halogens do not react with water, but we can look at their reaction with an alkali metal to find the pattern of reactivity for the family. – F2 + 2Na -> 2NaF explodes –Cl2 + 2Na -> 2NaCl vigorous –Br2 + 2Na -> 2NaBr slower – I2 + 2Na -> 2NaI have to heat to get

reaction going So what pattern is in the reactivity as you

move down the group?

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As you move down the group:

• Boiling/Melting point increases

• Reactivity decreases

• Density increases

Group 8A or 18: Noble Gases

• Do not react easily with anything, due stable electron configuration

• All other elements strive to reach noble gas configuration for maximum

stability by reacting with

other elements.

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Group 8A or 18: Noble Gases

As you move down the period:

• Reactivity decreases

• Melting point increases

• Boiling point increases

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Transition Metals: Group 3B-12B • Have many electrons that they are able to share, this allows them

to have many different ions: Cr2+, Cr3+, Cr5+, Cr6+

• Always lose electrons to make positive ions

• These varying charges gives variety of

colors to same element

• Malleable and ductile

• Conduct electricity and Shiny

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Charges of all the Families • Remember atoms that lose or gain electrons form

ions, and these are the charges each family forms.

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+1

+2 +3 +4 -3 -2 -1

0

Transition Metals

A Few More General Periodic Trends

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• As you move down a family atom size increases

• As you move left to right across the table

atom size decreases

• Where are the largest

atoms located?

3.9 Natural States of the Elements

Objectives: To learn the natures of some common elements.

Who is a solid, liquid or gas? • When we look at the elements on the periodic

table, who is a solid, liquid or gas in their natural state?

• Most elements are not found in their elemental state, most elements are found in compounds with other elements.

• Most elements on the periodic table are solids, so we will point out those who are gas or liquid?

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Liquids

• Only 2 elements in their elemental form are a liquid at 25 degrees Celsius: Mercury and Bromine

• Gallium and Cesium almost qualify, but they are solids 25 degrees Celsius, but melt around 30 degrees Celsius.

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Gases

• More elements exist in their elemental form as a gas, but there are some important distinctions to make about these gases.

• The noble gases are a gas, called monatomic gas. This means that the prefix mono- means one. And monatomic gases exist as individual atoms.

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Figure 3.13: A collection of argon

atoms.

Gases

• There is another group of gases called diatomic gases. The prefix di- means two. These elements travel in pairs as molecules.

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Figure 3.14: Oxygen gas

contains OXO molecules.

Figure 3.14: Nitrogen gas

contains NXN molecules.

Gases

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There are 7 elements that exist as diatomic molecules, you

will simply need to find a way to memorize these.

If you notice, all of the halogens fall in this category, and

then hydrogen, nitrogen, and oxygen.

You will also notice that 2 of these are not gases, make sure

you do not for get to include these in your diatomic list.

Noble Metals

• There is one group of metals which are relatively unreactive and thus called the noble metal. This group includes gold, silver and platinum.

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Periodic Table Diatomic Solids Solids

Diatomic Liquids Synthetically Made

Diatomic Gases

Monatomic Gases

Liquids

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3.10 Ions Objectives:

To describe the formation of ions from their parent atoms and learn to name them.

To predict which ion a given element forms by using the periodic table.

What is an ion?

• When we discussed atoms before, we were always looking at a neutral atom.

Neutral atoms always have equal numbers of protons and electrons.

protons = +1 charge

electrons = -1 charge

• When atoms have unequal numbers of protons and electrons, then the atom is a charged particle called an ion.

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Ion Facts • Ions are atoms, or groups of atoms, with a

charge.

• The charge is created by different numbers of protons and electrons.

• In an ion ONLY electrons can move.

• Atoms gain or lose electrons to become ions.

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Cations and Anions • There are 2 types of ions: cations and anions.

• Cations are ions with a positive charge. To form a cation, an atom has lost electrons.

Example: Na loses an electron and becomes Na+

• Anions are ions with a negative charge. To form an anion, an atom has gained electrons.

Example: Cl gains an electron and becomes Cl-

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Basic Names for Ions • Cations do not change names from their neutral

atoms.

Example: Magnesium loses 2 electrons and becomes Mg2+ which is named magnesium ion.

• Anions change the end of their name to –ide.

Example: Chlorine gains an electron and becomes Cl-. We would change the name from chlorine to chloride.

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Some Common Anion Names

• What would the names of the following ions be?

• Chlorine =

• Fluorine =

• Bromine =

• Iodine =

• Oxygen =

• Sulfur =

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How to Determine the Charge • When determining the charge for an atom we

can use the periodic table to help.

• The number of valence electrons determines the charge.

• All atoms want 8 valence electrons.

• If an atom has 1-3 valence electrons the atom will lose them to become positive.

• If an atom has 6-8 valence electrons the atom will gain electrons to become negative.

• We can determine the charge by looking at the periodic table.

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Figure 3.19: The ions formed by selected members

of groups 1, 2, 3, 6, and 7.

Practice • Determine the name and charge of the

following ions:

Potassium

Bromine

Calcium

Sulfur

Aluminum

Strontium

Cesium

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3.11 Compounds that Contain Ions

Objective:

To describe how ions combine to form neutral compounds.

Ions and Compounds

• Any time a bond is formed between 2 or more ions, we call this an ionic compound.

• In an ionic compound the overall charge must be zero.

• This means there must be cations and anions present.

Example: Na+ + Cl- -> NaCl

Each atom has a charge of +1 or -1, when we add this together it comes out to be zero.

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• If we add two ions together and they are not equal, then we must add another ion to balance the charge.

Example: Mg2+ + 2Cl- -> MgCl2

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Elements, Atoms, and Ions€¦ · 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types

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