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3-1
Chapter 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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Table 3.1
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are made up, mainly, of oxygen, carbon, hydrogen
and nitrogen.
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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.
• Some elements found in the human body are: 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 Not 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).
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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.
– John Dalton attempted to explain these observations in 1808.
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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.
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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. 3-20
Figure 3.6: Results of foil experiment if Plum Pudding model had been
correct.
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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. 3-23
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 subatomic 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.
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Section 3.7: Isotopes • .
• isotopes; two atoms of the same element (same number of protons) with different numbers of neutrons.
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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 X3-30
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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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
– 8 atoms have a mass of 23
– 2 atoms have a mass of 19
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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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 35 -> .7577x35 = 26.52
24.23% of mass is 37-> .2423x37= 8.96
Add the 2 together to get the atomic mass:
26.52+ 8.96 = 35.48
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Practice
• Oxygen has 3 isotopes 16O, 17O, 18O
99.76% of mass is 16O
0.04% of mass is 17O
0.20% of mass is 18O
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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Figure 3.12: Elements classified as metals and nonmetals.
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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
3-57
Group 7A or 17: Halogens or Hydrogen Family
• 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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Group 7A or 17: Halogens or Hydrogen Family
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
Figure 3.19: The ions formed by selected members of groups 1, 2, 3, 6, and 7.
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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.
3-70 Figure 3.14: Oxygen gas
contains OXO molecules.
Figure 3.14: Nitrogen gas
contains NXN molecules.
Diatomic molecules
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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.
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.
3-74
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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• IONS
• Cations lose e- and become positively charged.
• They keep their name.
• Anions gain e- and become negatively charged.
• Change the ending to ide.
• When combined to form a compound, cationscome first then anion.
• Net charge on the compound is 0.
• Ex. Magnesium oxide.
• What are the ions and what does the compound look like?
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