Chemical nomenclature are rules used when making names for chemical compounds. IUPAC, the International Union of Pure and Applied Chemistry, created and developed the nomenclature used most frequently worldwide. Their system of nomenclature names such chemicals as; ions, polyatomic ions, hydrates, molecular compounds, and acids and bases.
CHEMICAL FORMULAS When writing chemical formulas be aware of the differences between ion and compound formulas.
ION FORMULAS Zinc has a 2+ beside it's chemical symbol; we call this a zinc ion and the 2+ represents its ion charge. When writing an ion formula for a metal all you have to do is add the element's (metal) name in front of the word ion.
Examples: An ion formula for sodium (Na+) would be: sodium ion
To name an ion formula for a non-metal add an -ide ending to the non-metal name and put it in front of the word ion.
Example: Fluorine (F-) would be: fluoride ion
COMPOUND FORMULAS When writing compound formulas take the chemical symbol of your first element and put it in front of the chemical symbol of your second element. Remember, if you are making an ionic compound put the chemical symbol of the metal first. When Barium and Chlorine make an ionic compound Chlorine will have a subscript of 2, this represents the number of Chlorine ions in the compound.
MULTIVALENT IONS Some elements can form more than one ion charge (multivalent ion). An example of this is iron, which can have an ion charge of 3+ or 2+. When unsure of which charge to use, remember that the top number on the periodic table is the more common charge and should be used unless instructed not to. IUPAC uses roman numerals in parenthesis to show the charge. Classical (old) systems use latin names of elements and suffix -ic (the larger charge) or -ous (the smaller charge).
Some other Chemical names to remember are:
Ferr - Iron Cupp - Copper Mercur - Mercury Stann - Tin Aunn - Gold Plumb - Lead Wolf - Tungsten
COMPLEX IONS Complex ions are large groups of atoms that stay together during a chemical reaction. Almost all complex ions are anions (negative ions).
Elements can even make compounds with polyatomic ions!
HYDRATES Some compounds can form lattices that bond to water molecules.
To name hydrates, first write the name of the chemical formula. Then, add a prefix indicating the number of water molecules (mono=1, di=2, tri=3, etc.). Finally, add "hydrate" after the prefix. Example:
Next class, we will learn how to name molecular compounds!
The atomic symbol represents the nucleus. In order to know how many electrons one will need to draw, one will need to determine how many valence electrons that atom has. Electrons are drawn as dots and placed in four orbitals each containing a maximum of 2e- around the atomic symbol. Each orbital gets 1e- before they get paired with the left overs electrons.
Examples:
LEWIS DIAGRAMS FOR COMPOUNDS AND IONS
In bonds that are covalent, meaning a bond between two non-metals, electrons are shared.
1. First one must find the number of valence electrons for "x" atom in the molecule.
2. Draw the atoms so that they will fill each orbital which means that one needs to place one in each orbital before one begins pairing up.
Example:
-CH4
The two blue-boxed diagrams are Carbon and Hydrogen by themselvesand the white boxed one is the compound:
DOUBLE AND TRIPLE BONDS
In some cases the only way a covalent bond can be drawn is if they share more than one electron.
Examples:
IONIC COMPOUNDS
In ionic compounds electrons transfer from one element to another. One must determine the number of valence electrons on the cation (Mr. Doktor has a good way of remembering that a cation is positive; think of a cat scratching out a plus sign). Then, one must move these to the anion. Write the charges outside of the brackets.
Examples:
Here are two helpful websites that also explain Lewis diagrams:
Today Mr. Doktor explained that many properties of the elements change in a predictable way as you move through the periodic table; what is otherwise known as periodic trends. There are seven important periodic trends:
REACTIVITY Reactivity is the tendency of a substance to undergo chemical changes in a system. Metals and non-metals show different trends. The most reactive metal is Francium. The most reactive non-metal is Fluorine. Note that non-metals are located on the right side of the periodic table. As you move down the metals in the table, the reactivity increases. However, for non-metals, the reactivity increases as you move up. For the transition metals, reactivity increases as you move outwards (left and right).
Mr. Doktor even demonstrated the reaction of sodium with water. Here's a video showing what happens when you mix alkali metals with water:
ION CHARGE The ion charges of different elements depend on their group (column).
MELTING POINT Melting point is the temperature at which a substance changes from a solid to a liquid form. Elements in the center of the periodic table have the highest melting point (Tungsten). Noble gases have the lowest melting points. Also, melting point increases (until the middle of the table) starting from left and moving right. Exceptions are Carbon, Silicon, and Boron.
ATOMIC RADIUS Atoms get larger going down a group, and smaller moving left to right across each period. For example, Helium has the smallest atomic radius, while Francium has the largest atomic radius. Electrons with a larger principle quantum number are found in orbitals that extend farther away from the nucleus. This makes the atomic radius larger. As well, atoms that have more positive charge in their nuclei exert a stronger pull on the electrons in a given principal quantum level. In other words, atoms are made smaller when a stronger attractive force shrinks the electrons' orbitals.
IONIZATION ENERGY An atom's ionization energy is the energy needed to remove one of its electrons. That is, the energy needed to completely remove an electron from an atom. Ionization energy increases going up and to the right. These trends are the exact opposite of the trends for atomic radius. This is because both an atom's size and its ionization energy depend on how strongly its electrons are attracted to the nucleus. All noble gases have high ionization energy. Helium has the highest ionization energy. Francium has the lowest ionization energy (it is easy to get rid of an electron).
ELECTRONEGATIVITY An atom's electronegativity reflects its ability to attract electrons in a chemical bond; how much atoms want to gain electrons. It has the same trend as ionization energy. For example, Fluorine is the most electronegative element (an electronegativity of 4.0). Cesium and Francium both have electronegativities of 0.7. When two atoms form a chemical bond, the atom with the higher electronegativity would more strongly attract the bond's electrons. Electronegativity has the same trends as ionization energy.
DENSITY Density is very similar to atomic mass. Although, we did not go into great detail with this periodic trend.
Here's a video to help with your understanding of period tables and trends:
Before we explain the concepts of isotopes and atoms, at the beginning of class we learned what a "ground state" was and what an "excited state" was. The ground state of an atom is the state where the atom's electrons occupy the lowest energy orbitals available. The excited state of an atom is a state with more energy. When an orbital is skipped, the atom is in an excited state.
Now, onto today's lesson on atoms and isotopes!
ATOMS We all know that the atomic number is the number of protons in a single atom of an element. It is usually located in the top left of the element's square on the Periodic table. This tells you the atomic number (top left), the symbol (centre), the ion charge (indicated by a plus or minus sign on other periodic tables), and the atomic mass (top right). The atomic mass is only the average of all the isotopes of an element.
But, how does one find the number of neutrons in a single atom of an element? Well, my friend, it's much easier than you think. You only need to subtract the atomic number from the atomic mass (unless it is given, we can round this to the nearest whole number).
Atomic Mass - Atomic Number = The Number of Neutrons
(P + N) - (P) = (N)
Here are two links that also explain this step in greater depth:
Not all atoms of the same element are identical. An isotope is a form of an atom with the same number of protons but a different number of neutrons. In other words, Isotopes have the same atomic number, but the thing that separates them from being an atom is that they have different atomic MASS numbers , which is a result of having different amounts of neutrons within their nucleus.
The most common isotope of hydrogen has no neutrons at all; there's also a hydrogen isotope with one neutron (deuterium) and two neutrons (tritium).
Examples:
isotope: 54Fe --> mass number: 54 --> atomic number: 26 --> number of neutrons: 28
isotope: 56Mn --> mass number: 56 --> atomic number: 25 --> number of neutrons: 31
isotope: 237Np --> mass number: 237 --> atomic number: 93 --> number of neutrons: 144
isotope: 14C --> mass number: 14 --> atomic number: 6 --> number of neutrons: 8
ISOTOPE NOTATION
The atomic number is written as a subscript on the left of the element symbol. The mass number is written as a superscript on the left of the element symbol. Finally, the ionic charge (if any) appears as a superscript on the right side of the element symbol.
Let's look at hydrogen:
You'll notice that the smallest isotope, hydrogen-1, has only one neutron. Deuterium and Tritium each have more neutrons. Also, the nucleus does not repel itself, because the neutrons act as spacers.
A Mass Spectrometer is a device that can be used to determine the relative abundance and the mass of the isotopes of elements. To deflect a beam of charged particles, this device uses a charged field. The lightest particles will bend the most. We can measure where the particle lands on the screen, and in what abundance. Here is a diagram of a mass spectrometer:
We know that there are "preferred" combinations of neutrons and protons, at which the forces holding nuclei together seem to balance best. In fact, light elements tend to have about as many neutrons as protons, while heavy elements apparently need more neutrons than protons in order to stick together.
Let's say we found the following information:
And we want to find the average atomic mass. The 5 peaks in the mass spectrum shows that there are 5 isotopes of zirconium - with relative isotopic masses of 90, 91, 92, 94 and 96. So, we know:
zirconium-90 51.5
zirconium-91 11.2
zirconium-92 17.1
zirconium-94 17.4
zirconium-96 2.8
(51.5 x 90) + (11.2 x 91) + (17.1 x 92) + (17.4 x 94) + (2.8 x 96) = 9131.8. Then, we take 9131.8 and divide by 100. The average mass of these 100 atoms would be 9131.8 / 100 = 91.3 (to 3 significant figures). Find zirconium in the periodic table, and 91.3 is approximately the relative atomic mass of zirconium. Yay!
Next class, period tables and trends!
A very perplexing and fascinating topic was covered today in Chemistry class.
WHAT IS QUANTUM THEORY? Quantum Theory is a theory describing the behavior and interactions of elementary particles or energy states based on the assumptions that energy is subdivided into discrete amounts and that matter possesses wave properties. Energy of electrons is quantized (they can only have certain amounts of energy). Electrons exhibit wavelike behavior. It is impossible to know the exact position as well as momentum of an electron at any given instant. The quantum-mechanical model of an atom includes all of these ideas. Let's begin, shall we?
Theelectron is a particle that must be in an orbital in the atom. It is acloud of negativechargeor wave function. The cloud is most dense where the probability of finding the electron is highest. Electron density is the density of an electron cloud. An atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found. In other words, orbitals are in 3D space where the electrons most probably are. The energy of an electron is in its vibrational modes - like notes on a guitar string.
Photons are produced when high energy modes change to lower energy modes.
Like the Bohr model, the energy of the electron increases as n (principal energy levels) increases from 1 to 2 to 3, and so forth. Unlike the model, however, each principal energy level is divided into one or more sublevels, and each sublevel consists of one or more orbitals. There are several different kinds of orbitals. Each orbital has a different fundamental shape. They are designated by the letters s (spherical), p (dumbbell shaped), d, and f. Each orbital holds a maximum of 2 electrons.
S ORBITALS
The s orbital is a spherically-shaped region describing where an electron can be found, within a certain degree of probability.There is one sub-orbital. Each orbital holds a maximum of 2 electrons. Thus, there are 2 electrons in total.
P ORBITALS The p orbital is a dumbbell-shaped region describing where an electron can be found, within a certain degree of probability. There are 3 sub-orbitals. Each orbital holds a maximum of 2 electrons. Thus, there are 6 electrons in total.
D ORBITALS The d orbital has 5 sub-orbitals. Each orbital holds a maximum of 2 electrons. Thus, there are 10 electrons in total.
F ORBITALS The f orbital has 7 sub-orbitals. Each orbital holds a maximum of 2 electrons. Thus, there are 14 electrons in total.
ELECTRON SPIN
Electrons behave as though they were spinning on their own axis. The Pauli exclusion principle states that each orbital in an atom can hold at most 2 electrons and that these electrons must have opposite spins.
The most stable, lowest energy state of the atom is called the ground state. This is the way atoms normally exist.
AUFBAU PRINCIPLE
"Aufbau" is a German word meaning "building up." This principle states that electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.
PAULI EXCLUSION PRINCIPLE
This principle states that an orbital can hold a maximum of 2 electrons. The 2 electrons must spin in opposite directions. The electrons are said to be paired when they have opposite spins, and occupy an orbital. An unpaired electron is a single electron present in an orbital.
HUND'S RULE
This rule states that electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.
EXAMPLE:
*The atomic number of carbon is 6. The first 2 electrons go into the 1s orbital with their spins paired. The next 2 go into the 2s orbital with their spins paired. The fifth and sixth electrons occupy two separate 2p orbitals.
ELECTRON CONFIGURATION
The above is an orbital diagram, which shows how electrons populate orbitals.An up-arrow symbolizes an electron spinning in a counterclockwise direction. A down-arrow symbolizes an electron spinning in a clockwise direction. The direction of the arrows is based on the magnetic properties of the spinning electron, and the first electron entering an orbital is usually represented by an up-arrow.
There is an even more compact form of showing the electron configuration.
The pattern is: Poe
P is the principle quantum number.o is the orbital type (s, p, d, or f). e, the superscript, is the number of electrons in the orbital. The sum of the superscripts equals the number of electrons in the atom. The diagram below can help to write this notation. Start from the top and go across, working your way down. For the most part, the principle quantum number is the period number.
Notice that in the above periodic table, the d-orbitals (yellow) are labelled 3 in the fourth row. This is because the 4s sublevel is actually lower in energy than the 3d. So, the d-orbitals are an exception.
EXAMPLES:
*How many and what type of electrons are in cesium?
Starting from the top, moving across, and working your way down from the periodic table, you should arrive at:
1s22s22p63s23p64s23d104p65s24d105p66s1
But this is rather long, so there is a shortcut method. Write the chemical symbol of the preceding noble gas. Then, add in the electrons that are remaining.You should arrive at:
[Xe]6s1
*Using the shortcut method, write the configuration for Barium.
[Xe]6s2
*What about an ion, like Ba2+?
[Xe]
So, we are able to conclude that the chemical species Ba2+, Xe, and I- are isoelectronic, meaning , having the same electron configuration.
Next time we will discuss isotopes and mass spectrometers!
Last class, we determined that the Bohr Model is a planetary model in which the negatively-charged electrons orbit a small, positively-charged nucleus similar to the planets orbiting the Sun. Electrons orbit the nucleus in orbits. The lowest energy is found in the smallest oribt. Radiation is absorbed or emitted when an electron moves from one orbit to another. We continued our discussion on representing Bohr's atomic theory today. Hurray! Before representing Bohr's theory, you need to know that all atoms are electrically neutral, which means that they have the same number of protons and electrons. You should also know the two ways to represent Bohr's theory; the energy level model and the Bohr model. Electrons occupy shells which are divided into orbitals:
-2 electrons in the first shell -8 electrons in the second shell -8 electrons in the third shell
The last two shells are octets, and all atoms want stable octets (a full number of electrons in the last shell). The fourth shell holds 18 electrons. An ion will have a charge that is positive (less electrons than protons) or negative (more electrons than protons).
You can build an atom yourself, from scratch, here:
Click to Run
ENERGY LEVEL MODEL Here is an energy level model for Argon:
If you look at the periodic table, you will notice that the mass of Argon is 39.95 or 40 (which is the number on top). Its atomic number (the number of protons in the nucleus) is 18 (which is the number on the bottom). These numbers, along with the chemical symbol, Ar, make up what is called atomic notation. Argon happens to be a noble gas, so its shells are full (this is why noble gases do not react easily). Because atoms are neutral, the number of electrons is the same as the number of protons (18 in this case). Keeping in mind that there are 2 electrons in the first shell, 8 electrons in the second, and 8 electrons in the third, we can demonstrate that there are 18 electrons in total by writing 2e, 8e, and 8e above the atomic notation. In chemistry the energy level model is used more because the Bohr model doesn't represent the shape of the orbitals (each pair of electrons) correctly and is therefore incorrect (we'll learn more about these shapes later on!).
BOHR MODEL Here is a bohr model for Argon:
We draw dots to represent the electrons. Notice how there are 2 electrons in the first shell, 8 electrons in the second, and 8 electrons in the third, making a total of 18 electrons. Thus, the atom is neutral. Also, notice that the electrons are drawn in pairs. We draw 1 electron in each orbital before we start pairing them up.
Each row in the periodic table corresponds to the number of shells in a Bohr diagram. For example, there are 3 shells in the bohr diagram of Argon, because Argon is in the 3rd row of the periodic table. The periodic table group (vertical column) determines the number of valence electrons (8; in this case, a stable octet). The number within the unit's place identifies how many valence electrons are contained within the elements listed under that particular column. Helium, (2 valence electrons) and Groups 3-12 (transition metals) are an exception to this rule.
Some atoms have charges (ions). In the example below there is a Fluorine atom which has a charge of -1 because it needs just one more valence electron (electrons in the outer most shell) to have a stable octet, and there is a Calcium atom that has a +2 charge because it needs to get rid of 2 valence electrons to have a stable octet.
We went into a lot more detail regarding the atomic theory of Niels Bohr. It was a very BOHRing (not BORing) class today!
NIELS BOHR
Niels Bohr was one of the most influential scientists of the 20th century. He was a Danish physicist who made fundamental contributions to understanding atomic structure and quantum mechanics, and received the Nobel Prize for Physics in 1922.
EVIDENCE
Mr. Bohr expanded on Rutherford's theory (that the atom consisted of a positively charged nucleus with negatively charged electrons orbiting around it). Rutherford's model was inherently unstable (protons and electrons should attract each other)!
LINE SPECTRA
Bohr noticed that when white light was passed through a prism, the visible spectrum of light was obtained:
In contrast to this, he also saw that a line spectrum (which contains distinct lines of a particular color) was obtained when light of a particular color (like from the hydrogen discharge tube) was passed through a prism.
Bohr based his model on the energy (light) emitted by different atoms. Note that matter emits light when it is heated:
We can make objects and gasses glow by heating them up in a flame, or by passing electricity through them. The spectroscope spreads out the colors of the light, and we can identify the elements by the bright lines we see in the spectroscope. Black body radiation is the thermal radiation that would be emitted by a blackbody (that is, an ideal object that would absorb all of the radiation incident on it without reflecting any radiation). The distribution of energy in such radiation depends solely on the temperature of the source.
Light travels as photons. A photon is a quantum of visible light or other form of electromagnetic radiation demonstrating both particle and wave properties. It has neither mass nor electric charge but possesses energy and momentum. The energy photons carry depends on their wavelength.
Each element gives off a specific colour of light (emission spectra). It is unique to each element. When hydrogen is excited in a gas discharge tube, it emits its well-known line spectrum:
This allowed Bohr to put forward his model of the atom, which is still used, with modifications, to describe the atom today. To explain this emission spectra, he suggested that electrons occupy shells or orbitals.
BOHR'S THEORY
Bohr proposed that electrons only occupy certain orbits or shells in an atom. Also, each orbit represents a definite energy for the electrons in it. When they absorb energy, they move to a higher orbital. As they fall from a higher orbital to a lower one, they release energy as a photon of light.
Unit 2 has begun! This time, we will explore atoms; the smallest particle of an element that retains the chemical identity of the element. Learning about atomic theory is crucial when learning chemistry. An atomic theory is the theory of the nature of matter, which states that matter is composed of discrete units (atoms). Our understanding of atoms has changed over the years:
FOUR ELEMENT THEORY (ARISTOTLE) When things catch fire, moisture is released, air can be felt coming up from it, and ashes show the earth that it contains.
This theory states that there are four elements that make up everything → Earth, Fire, Wind, and Water. It simplified our view of the world, but can these types of matter be broken down even further?
DEMOCRITUS’ THEORY
All matter is composed of tiny particles. Democritus said that these particles are indivisible. This is when the idea of atoms were first mentioned.
However, this is only a conceptual model. There is no mention of a nucleus, or its constituents. Aristotle and the Greek philosophers did not accept Democritus’ theory. What holds these particles together? Why doesn’t matter fall apart?
LAVOISIER’S THEORY The Law of Conservation of Mass states that the mass of an isolated system remains the same. Therefore matter cannot be created or destroyed.
PROUST’S THEORY The Law of Constant Composition states that a chemical compound alwayscontains exactly the same proportion of elements by mass. For example, the mass of water is always 88.9% oxygen and 11.1% hydrogen.
DALTON’S THEORY Dalton concluded that the properties of matter could be explained in terms of atoms:
Each element is composed of extremely small particles called atoms
All atoms of a given element are identical, but they differ from those of any other element
Atoms are neither created nor destroyed in any chemical reaction
A given compound always has the same relative numbers and kinds of atoms
Atoms are solid, indestructible spheres (and look kind of like billiard balls). Different spheres exist for different elements. His theory is also based on the law of conservation of mass.
This theory does not mention subatomic particles, however.
J. J. THOMSON’S THEORY He created the raisin bun model, and said that solid positive spheres have negative particles in them. He proposed the first atomic theory that includes protons (positive),and electrons (negative). Thomson demonstrated that electrons existed with a cathode ray tube – a vacuum tube in which a stream of electrons is produced and directed onto a fluorescent screen (e.g. in a television or visual display unit, creating images and text).
RUTHERFORD’S THEORY He showed that all atoms have a positive, dense center with electrons outside of it. The famous gold foil experiment:
He also suggested that atoms are mostly empty space and he explains why electrons spin around the nucleus. His model is a planetary model.
BOHR’S THEORY
The advent of quantum mechanics. He applied quantum theory to Rutherford's atomic structure by assuming that electrons travel in stationary orbits defined by their angular momentum. We’ll learn more about Bohr later on in the unit.
