Chemestry chapter 1 and 2
SECTION 1.2 – Atomic Structure
- understand the difference between empirical and theoretical knowledge.
- THEORIES are dynamic, they continually undergo refinement and change. They must:
- describe observations in terms of non-observable ideas
- explain observations by means of ideas or models
- Successfully predict results of future experiments
- be as simple as possible in concept and application
Early Theories
1) Democritus
- all matter could be divided into smaller and smaller pieces until a single indivisible particle is reached. This is the ATOM.
- believed that different atoms were of different sizes and geometric shapes and are in constant motion.
- believed that there is empty space between atoms.
2) Dalton
- developed the atomic theory of matter (learned in grade nine)
Include the 4 Postulates if needed
- Said that atoms are neither created nor destroyed in a chemical reaction and since atoms are indivisible, they must simply be rearranged. This way we end up with the same number of atoms at the end of a reaction as we had before the reaction began. This is called THE LAW OF CONSERVATION OF MASS.
- Since atoms of an element have identical properties such as mass and combine in constant ratios, every compound must have a fixed, definite composition. This is THE LAW OF CONSTANT COMPOSITION.
3) Thomson
- blueberry muffin experiment which discovered the electron.
- Electrons are negatively charged and are a just a small part of the atom.
- They are spaced evenly in a positively charged sphere.
4) Rutherford
- shot small positively charged particles at a piece of gold foil.
- He predicted they would all go through but some were deflect away which suggested that some type of positive charge existed in the centre of the atom. This is later found to be the nucleus.
- He coined the term proton for the smallest unit of positive charge in an atom.
5) Chadwick
- found that the atomic nucleus must also contain heavy neutral particles which he called neutrons.
6) Bohr
- He experimented with applying electricity and thermal energy to hydrogen gas.
- Found that the H atoms emitted light when they were “excited” by the additional energy.
This leads us to the idea that electrons orbit the nucleus in definite energy levels.
Each orbit has a specific quantity of energy assigned to it and therefore the orbits are often called energy levels.
As an electrons falls back from a higher energy level, it emits light energy. Where it started and where it falls determines the colour of the light.
In this model, each energy level can only hold a certain number of electrons. The arrangement is as follows:o
1st level – 2
o
2nd level – 8
o
3rd level – 18
Bohr-Rutherford Diagrams
This representation of the atom include a nucleus, with the number of protons and neutrons, the electrons all placed in their appropriate orbitals and what is called the valence electrons. These are the electrons that exist in the outermost energy level.
Atomic Mass & Atomic Number
- The number of protons in the nucleus of an atom determines the identity of the atom and is referred to as the ATOMIC NUMBER. It is symbolized by the letter (Z).
- Atomic number and nuclear charge are the same for all atoms.
- Since atoms are electrically neutral, the atomic number also must represent the number of electrons present in an atom.
- The sum of the number of nuclear particles in an atom is known as the ATOMIC MASS and is represented by the letter (A). Therefore, the number of neutron in an atom can be found by the formula:
N = A – Z
As an example of how small an atom is, visualize the head of a pin. There are about 8.0 x 1019 (80 000 000 000 000 000 000 atoms).
Because this number is so large we have determined that the standard atomic mass unit is equal to 1/12 the mass of a carbon-12 atom.
The masses of all other atoms are therefore measured in relation to the mass of carbon-12.
Example: F-19 has a mass that is 19/12 of that of a carbon-12 atom, and each H atom has a mass of 1/12 of the C-12 atom.
Therefore the mass of a C-12 atom is 1u.
Practice Questions
p.16 # 1-5, 10
Section 1.3 – Ions & The Octet Rule
The Octet Rule
When atoms have a full valence shell they have a special stability (which is desired by other atoms).
When considering the first 18 elements, a full valence shell would have 8 electrons. He, Ne and Ar all have full valence shells (2, 8 and 8 respectively).
Atoms of the other elements do not have full valence shells but when they combine with other atoms they tend to achieve this state. This is known as the octet rule.
There are only three possible ways for an atom to achieve that full valence shell. They can either gain, lose or share electrons.
When and atoms either gains or loses electrons, it becomes an ion (either positively or negatively charged).
Cations – Positive Ions
The metals, which exist to the left of the staircase on the PT, tend to have very few valence electrons.
Because of this they tend to lose e- in order to attain a stable octet.
When e- are lost, there are now more + charges in the atoms then there are – charges so the atom now has an overall positive charge.
Example: Sodium
Negative Ions – Anions
Elements on the right side of the PT have almost full valence shells, and therefore tend to gain e- in order to obtain the full octet.
When an atom gains e-, it now has more – charge that + charge (protons) and therefore has an overall negative charge.
Example: Chlorine
In either case when an ion is formed the chemical symbol will change.
Example: Na becomes Na+1 and Cl becomes Cl-1
If more than one e- is gained or lost, the charge on the corresponding new ion will reflect the number of e- in question.
Example: Mg becomes Mg+2 (as it loses 2e-).
Multivalent Ions
The metals is the middle of the PT are referred to as the transition metals. Many of them have the unique ability to form multiple ions.
Example: Copper can form both Cu+1 and Cu+2 ions.
Polyatomic Ions
An ion that contains more than just one atom is called a polyatomic ion.
Most of these type of ions are composed of oxygen, sulfur, phosphorus, chlorine and carbon. Most of them are also anions.
Just like simple ions, their charge represents their desire to achieve a stable octet.
Below are two charts that show a)ions in the human body and b)Common polyatomic Ions
Section 1.4 – Isotopes, Radioisotopes and Atomic Mass
ISOTOPES
- They are forms of elements in which the atoms have the same number of protons but a different number of neutrons. In other words, they have the same atomic number (Z) but a different atomic mass (A).
- It is the protons and electrons that are responsible for the chemical behavior of an atom. This means that isotopes share chemical properties even though their masses are slightly different. The physical properties of these isotopes however can be drastically different.
- On this scale, the protons and neutrons have a mass of about 1u while electrons have a mass of about 0.000 55u
- Naturally occurring magnesium is a combination of three isotopes. Mg-24, Mg-25 and Mg-26.
- An average sample would contain 78.7% Mg-24, 10.1% Mg-25 and 11.2% Mg-26. This is known as the isotopic abundance.
- These values are determined using a mass spectrometer.
RADIOISOTOPES
- These are atoms of unstable isotopes which decay and emit radiation as the nucleus changes.
- These isotopes are said to be radioactive.
- These isotopes give off one of three types of radiation: alpha, beta or gamma rays. The table below summarizes the characteristics of these three type of radiation.
- Each radioisotope has a property called HALF-LIFE. This is the time taken for half of the original number of radioactive atoms to decay. The half-lives of different radioisotopes vary dramatically.
Determining Atomic Mass
- Consider the following problem of finding the average height of a group of 8 people.
- The atomic mass of magnesium is 24.3u. How is this determined?
- We need to take the weighted average of the masses of all the isotopes of magnesium in order to get this number.
Section 1.5 – The Periodic Table & Periodic Law
- In the modern periodic table, elements are arranged with similar chemical properties in vertical columns called GROUPS.
- Elements whose properties change from metallic to non-metallic as we move across a horizontal row and called PERIODS.
- Some of the most important groups in the periodic table are indicated below.
- All elements are arranged in order of increasing atomic number. This allows us to observe trends as we move across periods and is known as The Periodic Law.
Alkali metals
- soft silver coloured elements
- solids at SATP
- exhibit metallic properties
- react violently with water to form basic solutions and give off hydrogen gas
- react with halogens to form compounds like sodium chloride
Alkaline earth metals
- very light, very reactive metals
- form oxide coatings when exposed to air
- react with oxygen to form compounds with formula MO
- react with hydrogen to form compounds with the formula XH2.
- react with water to liberate hydrogen.
Noble Gases
- gases at SATP
- low melting and boiling points
- very unreactive
Halogens
- can be solid, liquid or gas at SATP
- exhibit non-metallic properties
- extremely reactive
Transition Metals
- exhibit a wide range of physical and chemical properties
- usually strong, hard metals with high melting points and good conductors of electricity
- variable reactivity
- form ions with variable charges
- One thing we notice in terms of Bohr-Rutherford diagrams, is that as we look down a group in the PT, the number of valence e- is the same for each element in a group.
- The number of shells increases by one as we move down the group but the valence e- remains the same.
- We can use Lewis Dot Diagrams as an alternative to BR diagrams. The Lewis diagrams only show the element symbol and the valence e-.
- Example: Nitrogen
- It is important to remember however all the knowledge that is assumed when using Lewis diagrams. The number of protons and neutrons as well as the remaining non valence e-.
Section 1.7 – Periodic Trends & Atomic Properties
- All electrons carry a certain amount of energy. The amount of energy determines what three dimensional pathway called an ORBIT that the electron is in.
- The shells are designated by the principal quantum number n which can be any integer from 1 to infinity. They are also sometimes represented by a letter starting with k,l,m...
- Bohr theorized that if an electron gains extra energy it could jump to a higher unfilled energy level (shell) farther away from the nucleus.
- He referred to this jump as a TRANSITION. The quantity of energy (quantum) required to cause a transition is equal to the difference in energy between the orbitals.
- When an electron falls back from a higher orbital it releases the amount of energy that was necessary to make it jump up in the first place.
- When electrons are in the lowest energy level that it can occupy it is said to be in its GROUND STATE.
THE QUANTUM MECHANICAL THEORY
- retains the concept that electrons fill successive shells, each of which is designated by the principal quantum number (n).
- describes electrons in atoms in terms of their energy and probability patterns (the probability of finding and electron of a specified energy within a specific region of space about the nucleus called THE ELECTRON CLOUD.
- To fully designate all probabilities, the energy levels are subdivided into an increasing number of subshells as the principal quantum number increases.
- Just as there is a maximum number of electrons for each shell there is also a maximum number for each subshell as well.
- According to Bohr, the maximum number of electrons that a given shell can hold is given by the equation:
2n2
n = principal quantum number
- According to this the first shell could hold 2 x (1)2 = 2
the second shell 2 x (2)2 = 8
the third shell 2 x (3)2 = 18
- The 17 electrons in a chlorine atom would be distributed over three shells. 2 in the first, eight in the second and seven in the third. These electrons (and any other) that occupy the outside shell are called the VALENCE ELECTRONS.
- Some rules to remember:
1) for groups 1 and 2 the number of valence electrons is the same as the group number.
2) in groups 13 to 18, the number of valence electrons is the same as the second digit of the group number.
Atomic Properties
- WITHIN ANY GIVEN PERIOD, chemical reactivity tends to be high in group1, lower in the middle and increase to a maximum in the group 17 elements.
- Metals react differently from non-metals. As we move down a group of metals, the chemical reactivity goes up. As we move down a group of non- metals, the chemical reactivity goes down.
Atomic Radius
- It is the distance from the centre of an atom (inside the nucleus) to the outermost electrons.
- Atomic radius decreases from left to right across each period.
- As we move across a period from group 1 to group 17 the number of protons in the nucleus increases by one, therefore so does the nuclear charge.
- As the nuclear charge increases the number of filled energy levels, electrons shielding the valence electrons form the nucleus remains the same.
- The attraction between the nucleus and the valence electrons is stronger as protons are added. This attraction pulls the electrons closer to the nucleus and decreases the size of the atom.
- In a given group, the atomic radius increases from top to bottom because of the increasing number of energy levels.
- Each extra level places electrons farther from the nucleus. This happens because the inner electrons which fill the inner shells are shielding the outer electrons from the pull of the nucleus.
- Each additional level reduces the attractive force of the nucleus and thereby leads to a larger atomic radius.
Ionic Radius
- The removal of an electron from an atom results in an ION.
- The positive ion has a smaller ionic radius than the atomic radius because there is now one less energy shell. The charge on this ion is positive because an electron is missing.
- Positive ions are always smaller than their neutral counterparts.
- The attraction of the nucleus for the electrons in positive ions increases across the period, and the ionic radius gets smaller.
- When an atom gains an electron it forms a negative ion and the ionic radius is now larger than that of the neutral atom.
Ionization Energy
- It is the amount of energy required to remove an electron from an atom or ion in the gaseous state.
- The amount is not constant but rather depends on which electron is being removed.
- The first ionization energy is the amount of energy required to remove the most weakly held electron from its neutral atom. This weakly held electron is always in the valence shell where it is most shielded from the attraction of the nucleus.
- This process is shown by: X(g) + energy ------> X+(g) + e-
- The second ionization energy is the same only it is to remove a second electron.
HOW ARE THE TWO RELATED?
- In general, the first ionization of metals is lower than those of non- metals.
- The IE decreases as you move down a group in the table. The decrease is related to the increase in size of the atoms from top to bottom in the same group.
- As the atomic radius increases, the distance between the valence electrons and the nucleus also increases. The result is a weaker attraction between the negatively charged electron and the positively charged nucleus. Therefore, less energy is needed to remove those electrons.
- The stronger the attraction between the nucleus and the electrons in question, the greater the amount of energy needed to remove those electrons.
ELECTRON AFFINITY
- When an electron is added to a neutral atom in the gaseous state, energy is usually released.
X(g) + e- -----> X-(g) + energy
- The energy associate with this ionization process is referred to as electron affinity.
- This process does not always result in the release of energy though. When the attractive force outweighs the repulsive force, energy is released as an electron joins the atom to form an ion. When the repulsive force is greater, the electron must be injected to form the ion.
- This is shown by second electron affinities. Whenever we add a second electron to an ion with a charge of negative one, the repulsive force is always greater than the attractive forces.
- Electron Affinity shows the same trends as ionization energy.
-It decreases as you move down a group.
-It increases as you move across a period.
ELECTRONEGATIVITY
- It is the ability of an atom to attract electrons. It combines ionization energies, electron affinities and some other measures of reactivity.
- Fluorine has been assigned a value of 4.0. It is the element considered to have the greatest ability to attract electrons. All other elements have their values based on that of fluorine.
Summary Diagram
SECTION 2.1 – Ionic Compounds
- Ionic compounds are those that are composed of a metal joined to a non-metal.
- Molecular compounds are those that are composed of two non- metals joined together.
- One of the ways we can distinguish between the two is to use electrical conductivity. Ionic compounds form solutions that conduct electricity while molecular compounds do not.
- Theses substances that form solutions that conduct electricity are called ELECTROLYTES.
IONIC BONDING
- When chemical bonds are formed, it is the valence electrons that are responsible for their formation.
- Ionic bonds are formed when one or more electrons is transferred form a metal to a non- metal, leaving the metal as a positive ion (cation) and the non-metal as a negative ion (anion). The two oppositely charged ions are attracted to each other and are held together by a force called an IONIC BOND.
- Ionic compounds are solids at SATP, with high melting points and they are electrolytes.
- Pure ones are electrically neutral, therefore the sum total of all the ions in it is zero.
- The most common arrangement of these ions is a regular three dimensional patterns called a CRYSTAL LATTICE.
- The structure of the lattice varies depending on the size and charges of the ions that make it up.
- The lattice requires a great deal of energy to break up because of the large electrostatic attraction of its pieces.
- Once the crystal is broken however, repulsion between like ions will force the formation of two separate crystals.
- When these crystals are dissolved in water the ions dissociate form one another.
Take NaCl for example: NaCl(s) ------> Na+(aq) + Cl-(aq)
- Elements in the same chemical group tend to participate in the same types of chemical reactions and from compounds (in this case ionic) with the same general formula.
- For example, group1 elements react with group17 elements to form ionic compound with the formal MX. They are referred to as metallic halides.
- Elements in group2 show a similar trend with group17 to produce metallic halides with the formula MX2.
- In general the addition of a group 1 or group2 metal to water will produce hydrogen gas and a basic ionic compound.
- Ions that have eight valence electrons resemble the noble gases and are therefore very unreactive. This is because they are said to have a STABLE OCTET.
- Each group in the periodic table will try and obtain a stable octet when forming ions. Group1 elements lose one electron to form a +1 ion while Group16 elements gain two electrons to form a -2 ion.
- Hydrogen is special because can either give up its only electron or it can gain one electron to be like helium.
- Electron dot diagrams which consist of the chemical symbol for the element plus dots representing the number of valence electrons.
- Sodium has one valence electron. By transferring it to another substance it will now have the same configuration as the nearest noble gas neon.
EXAMPLE:
- Chlorine has seven valence electrons. By gaining one from another source it will have the same configuration as its nearest noble gas argon.
EXAMPLE:
- Another way of describing the formation of ionic compounds and bonds is the difference in electronegativity between two atoms.
- When atoms of low electronegativity are in close proximity to those of high electronegativity, one or more electron is transferred from the atom with low electro.
- Two oppositely charged ions are formed and the attraction is there. This forms our crystal lattice.
SECTION 2.2 – Molecular Elements & Compounds
- The two compounds acetylene (C2H2) and benzene (C6H6) have the simplest ratio formula CH, but otherwise they are much different.
- A simplest ratio formula indicate only the relative number of atoms in a molecular compound. They do not give any information about the ACTUAL number of atoms present.
- Molecules that consist of only two atoms of one element are called DIATOMIC, while those with more than two are called POLYATOMIC.
- Recall hydrogen. It can obtain a stable octet by either gaining or losing one electron. Two separate hydrogen atoms however can both attain a stable octet if they share their one electron with the other. When this sharing takes place a COVALENT BOND is formed.
- In Electron dot diagrams the shared pair of electrons is replaced by a single dash to represent the covalent bond.
EXAMPLE:
- When the electrons are shared, each of the two atoms has a stable octet. This is referred to as the OCTET RULE.
- When the non-bonded pairs of electron are not shown in the Lewis dot diagram it is referred to as a STRUCTURAL FORMULA.
- In general, elements that need only one electron to fill their outer shell tend to form single bonds.
- The number of covalent bonds (shared pairs) that an atom can form is known as the atoms’ BONDING CAPACITY.
- When we draw structural formulas of molecules with multiple atoms, the central atom is the one with the highest bonding capacity. Electronegativity is another way to decide. The atom that has the lowest electronegativity becomes the central atom.
Examples: HCl, H2O and CO2
Drawing Lewis Structure for Polyatomic Ions
Example: ClO3-
1) Arrange the symbols of the elements so that one atom is the central element. Spread the other atoms around the central one.
2) Count all the valence e- in the molecule. Record that number.
3) Place two e- (a single covalent bond) between the central atom and each of the surrounding atoms.
4) Place lone pairs around each of the surrounding atoms (except hydrogen) to satisfy the octet rule.
5) Determine how many e- remain to be placed.
6) Place those remaining e- on the central atom in pairs.
7) If the central atom does not have a full octet, move lone pairs from surrounding atoms into bonding positions (forming either double or triple bonds).
8) If the surrounding atoms have complete octets and there are e- remaining, add these e- as lone pairs onto the central atom.
9) Replace shred pairs with bond dashes and write the charge outside a square bracket.
Practice Questions
p.69 #2-4, 6
SECTION 2.3 – Chemical Bonding & Electronegativity
-We have previously looked at the concept of electronegativity, or the ability of an atom to attract e-.
We can use this concept to determine whether a compound, when formed, will be ionic or molecular.
We use the difference in electronegativity between the atoms to make the distinction. If the difference is greater than 1.7, the compound will be ionic, below that value, it will be molecular.
POLAR COVALENT BONDS
- As we have seen, when there is sharing of electrons between atoms a covalent bond is formed. The sharing is exactly EQUAL when the two atoms are identical. Such as in the chlorine to chlorine bond.
- However, when a molecule like hydrogen chloride exists, where the electrons are shared between different atoms, the sharing is UNEQUAL. The bonding electrons spend more time near one atom than the other.
- In the H-Cl bond, the electrons spend more time near the chlorine than the hydrogen because of chlorine’s greater attraction for electrons.
- Due to this uneven sharing, the hydrogen is slightly positive and the chlorine is slightly negative.
- Although chlorine attracts the electrons more strongly than hydrogen, the attraction is not strong enough to cause a transfer as in an ionic bond. The bond is somewhere between an ionic bond and a covalent bond and is called a POLAR COVALENT BOND.
- The element with the greatest electronegativity will always have the slight negative charge.
EXAMPLE:
- The difference in electronegativity of two bonded atoms is what determines the type of bond that forms. It also provides a measure of the polarity of the bond. The greater the difference, the more polar the bond.
- By convention, a difference in electronegativity greater than 1.7 indicates an ionic bond. There are some compounds that have ionic properties whose electronegativity is less than 1.7.
Practice Questions
#1, 2, 5, 6, 7
SECTION 2.4 - NAMES AND FORMULAS OF COMPOUNDS – Chemical Nomenclature
- The system in place today for chemical nomenclature is developed by the International Union of Pure and Applied Chemistry (IUPAC).
- We will examine how to move from the chemical formula to the chemical name and vice versa.
BINARY IONIC COMPOUNDS
- They are simplest compounds found. They consist of two types of monatomic ions (ions consisting of one charged atom).
- For these, the metal cation is always written first, followed by the non-metal anion.
For Naming the Compound
- The name of the metal is stated in full, and the name of the non-metal ion has its suffix changed to “IDE”.
EXAMPLE:
For Writing the Formula
- Write the chemical symbol for each element involved.
- Place the valence on each of them (The charge on an ion is called its VALENCE).
- Cross the valences down to become subscripts and drop out the charges. (The cross over method).
- EXAMPLE:
- Most transition metals and some representative metals can form more than one kind of ion. Iron for example can have a charge of +2 or +3. Elements such as this are said to be MULTIVALENT.
For Writing the Name
- To name these compounds (with multivalent ions), we write the name of the metal in full and include the charge of the ion in Roman numerals in brackets. The non-metal suffix is again IDE.
EXAMPLE:
- In order to ensure the charge on the multivalent ion is correct we always check the non-metal ion first and then make adjustments if necessary.
EXAMPLE:
For writing the Chemical Formula
- Take the symbols for each element and then add the charges (valences) to them. The valence for the multivalent metal will be indicated in a bracket with Roman numerals. Then use the crossover method.
Example:
PRACTICE QUESTIONS
p.75 #1&2 p.77 #1
COMPOUNDS WITH POLYATOMIC IONS
- Compounds such as this are classified as tertiary compounds. They have a metal ion and a polyatomic ion (a covalently bonded group of atoms possessing a net charge).
- Those polyatomic ions that include oxygen are called OXYANIONS.
- To determine the name of these compounds, the first part is the name of the metal cation in full. The second part is more difficult.
Chlorine and oxygen can form 4 different oxyanions.
Hypochlorite ClO-
Chlorite ClO2-
Chlorate ClO3-
Perchlorate ClO4-
- For each of the four, the root is CHLOR. The suffixes and prefixes vary according to the number of oxygen atoms according to the rules below.
1) The PER-ATE oxyanion has one more oxygen than the ATE anion.
2) The ITE ion has one oxygen less than the ATE ion.
3) The HYPO-ITE ion has two oxygens less than the ATE ion.
- The following list are the ‘root’ polyatomic ions that you will need to memorize going forward.
Chlorate ClO3-
Carbonate CO3-2
Nitrate NO3-
Hydroxide OH-
Phosphate PO4-3
Sulfate SO4-2
Ammonium NH4+
HYDRATES
- These are tertiary compounds that contain water in the crystal structure.
- When heat is applied to a hydrate, it will form water vapour and an associated ionic compound. This indicated how loosely the water is held in the compound.
- When the water is removed the product is now ANHYDROUS.
- IUPAC names for ionic hydrates also indicate the number of water molecules by a Greek prefix. So Copper (II) Sulfate Pentahydrate is written as CuSO4 5H2O .
NAMING MOLECULAR COMPOUNDS
- These are compounds composed of two non-metals. There is a wide variety of these because the same two non-metals may combine in different proportions to form different compounds.
- In naming these types of compounds, a Greek prefix is again used. This time to indicate the name of each element in the compound. The prefix indicates the number of molecule of each element in the molecule.
- The prefixes are found here:
1- Mono 6 – Hexa
2- Di 7 - Hepta
3 –Tri 8 - Octa
4 – Tetra 9 - Nona
5 – Penta 10 - Deca
Practice Questions
p. 81 #1, 3-7 Plus extra practice worksheets (on request)
Mixed Naming Worksheet #2 W 303 Everett Community College Tutoring Center Student Support Services Program
Write the names of the following chemical compounds:
1) Fe(CN)3 ______________________________________
2) K4C ______________________________________
3) AuF3 ______________________________________
4) N2O______________________________________
5) Ag3N ______________________________________
6) CF4 ______________________________________
7) MgI2 ______________________________________
8) NiO2 ______________________________________
9) P2S5 ______________________________________
10) SnSe2 ______________________________________
Write the formulas of the following chemical compounds:
11) tricarbon octafluoride ____________________________________
12) lithium acetate ______________________________________
13) iron (II) arsenide ______________________________________
14) titanium (IV) acetate ______________________________________
15) gallium sulfide ______________________________________
16) ammonium carbide ______________________________________
17) ruthenium (II) nitrate ______________________________________
18) copper (I) oxide ______________________________________
19) potassium hydroxide ______________________________________
20) sodium phosphate ______________________________________
Write the names of the following chemical compounds:
1) Fe(CN)3 iron (III) cyanide
2) K4C potassium carbide
3) AuF3 gold (III) fluoride
4) N2O dinitrogen monoxide
5) Ag3N silver nitride
6) CF4 carbon tetrafluoride
7) MgI2 magnesium iodide
8) NiO2 nickel (IV) oxide
9) P2S5 diphosphorus pentasulfide
10) SnSe2 tin (IV) selenide
Write the formulas of the following chemical compounds:
11) tricarbon octafluoride C3F8
12) lithium acetate LiC2H3O2
13) iron (II) arsenide Fe3As2
14) titanium (IV) acetate Ti(C2H3O2)4
15) gallium sulfide Ga2S3
16) ammonium carbide (NH4)4C
17) ruthenium (II) nitrate Ru(NO3)2
18) copper (I) oxide Cu2O
19) potassium hydroxide KOH
20) sodium phosphate Na3PO4