The Arrangement of Atoms in the Periodic Table

1 The Arrangement of Atoms in the Periodic Table [Material is from The Free High School Science Texts: Textbooks for High School Students Studying the Sciences Chemistry Grades 10 12. www.fhsst.org unless otherwise noted. All graphics are copyright free.] The periodic table of the elements is a tabular method of showing the chemical elements. Most of the work that was done to arrive at the periodic table that we know, can be attributed to a man called Dmitri Mendeleev in 1869. Mendeleev was a Russian chemist who designed the table in such a way that recurring ( periodic ) trends in the properties of the elements could be shown. Using the trends he observed, he even left gaps for those elements that he thought were missing. He even predicted the properties that he thought the missing elements would have when they were discovered. Many of these elements were indeed discovered and Mendeleev s predictions were proved to be correct. To show the recurring properties that he had observed, Mendeleev began new rows in his table so that elements with similar properties were in the same vertical columns, called groups. Each row was referred to as a period. One important feature to note in the periodic table is that all the non-metals are to the right of the zig-zag line drawn under the element boron. The rest of the elements are metals, with the exception of hydrogen which occurs in the first block of the table despite being a nonmetal. One change since it was first invented: the elements in the modern table are lined up by the number of protons (atomic number) each has. Groups in the periodic table A group is a vertical column in the periodic table, and is considered to be the most important way of classifying the elements. If you look at a periodic table, you will see the groups numbered at the top of each column. The groups are numbered from left to right as follows: 1, 2, then an open space which contains the transition elements, followed by groups 3 to 8. These numbers are normally represented using roman numerals. In some periodic tables, all the groups are numbered from left to right from number 1 to

2 number 18. In some groups, the elements display very similar chemical properties, and the groups are even given separate names to identify them. The characteristics of each group are mostly determined by the electron configuration of the atoms of the element. The number of valence electrons of an element corresponds to its group number on the periodic table. Group 1: These elements are known as the alkali metals and they are very reactive. Group 2: These elements are known as the alkali earth metals. Each element only has two valence electrons and so in chemical reactions, the group 2 elements tend to lose these electrons so that the energy shells are complete. These elements are less reactive than those in group 1 because it is more difficult to lose two electrons than it is to lose one. Group 3 elements have three valence electrons. Group 7: These elements are known as the halogens. Each element is missing just one electron from its outer energy shell. These elements tend to gain electrons to fill this shell, rather than losing them. Group 8: These elements are the noble gases. All of the energy shells of the halogens are full, and so these elements are very unreactive. Transition metals: The differences between groups in the transition metals are not usually dramatic. It is worth noting that in each of the groups described above, the atomic diameter of the elements increases as you move down the group. This is because, while the number of valence electrons is the same in each element, the number of core electrons increases as one moves down the group. Periods in the periodic table A period is a horizontal row in the periodic table of the elements. Some of the trends that can be observed within a period are highlighted below: As you move from one group to the next within a period, the number of valence electrons increases by one each time. In general, the diameter of atoms decreases as one moves from left to right across a period. Consider the attractive force between the positively charged nucleus and the negatively charged electrons in an atom. As you move across a period, the number of protons in each atom increases. The number of electrons also increases, but these electrons will still be in the same energy shell. As the number of protons increases, the force of attraction between the

3 nucleus and the electrons will increase and the atomic diameter will decrease. Ionization energy (energy needed to remove an e - from the atom) increases as one moves from left to right across a period. As the valence electron shell moves closer to being full, it becomes more difficult to remove electrons. The opposite is true when you move down a group in the table because more energy shells are being added. The electrons that are closer to the nucleus shield the outer electrons from the attractive force of the positive nucleus. Because these electrons are not being held to the nucleus as strongly, it is easier for them to be removed and the ionization energy decreases. In general, the reactivity of the elements decreases from left to right across a period. Metals, non-metals, metalloids The elements in the Periodic Table can also be divided according to whether they are metals, semi-metals or non-metals. On the right hand side of the Periodic Table is a dark zigzag line. This line separates all the elements that are metals from those that are non-metals. Metals are found on the left of the line, and non-metals are those on the right. Metals, semi-metals and nonmetals all have their own specific properties. Metals Examples of metals include copper (Cu), zinc (Zn), gold (Au) and silver (Ag). On the Periodic Table, the metals are on the left of the zig-zag line. There are

4 a large number of elements that are metals. The following are some of the properties of metals: Thermal conductors Metals are good conductors of heat and are therefore used in cooking utensils such as pots and pans. Electrical conductors Metals are good conductors of electricity, and are therefore used in electrical conducting wires. Shiny metallic luster Metals have a characteristic shiny appearance and are often used to make jewellery. Malleable This means that they can be bent into shape without breaking. Ductile Metals can stretched into thin wires such as copper, which can then be used to conduct electricity. Melting point Metals usually have a high melting point and can therefore be used to make cooking pots and other equipment that needs to become very hot, without being damaged. Non-metals In contrast to metals, non-metals are poor thermal conductors, good electrical insulators (meaning that they do not conduct electrical charge) and are neither malleable nor ductile. The non-metals are found on the right hand side of the Periodic Table, and include elements such as sulfur (S), phosphorus (P), nitrogen (N) and oxygen (O). Metalloids (also called semi-metals) Semi-metals have mostly non-metallic properties, but tend to fall between metals and non-metals in most properties. For instance, they don t conduct electricity as well as do metals, but they do allow some electricity to flow through them. One of their distinguishing characteristics is that their conductivity increases as their temperature increases. This is the opposite of what happens in metals. The semi-metals include elements such as silicon (Si) and germanium (Ge). Notice where these elements are positioned in the Periodic Table. B and Al are not metalloids; all the other elements which share a full side with the heavy black line on the periodic table are metalloids.

Formation of Ions A change in the number of electrons of an atom does not change the type of atom that it is. However, the charge of the atom will change. If electrons are added, then the atom will become more negative. If electrons are taken away, then the atom will become more positive. The atom that is formed in either of these cases is called an ion. Put simply, an ion is a charged atom. A positively charged ion is called a cation e.g. Na +, and a negatively charged ion is called an anion e.g. F. The charge on an ion depends on the number of electrons that have been lost or gained. Look at the following examples. Notice the number of valence electrons in the neutral atom, the number of electrons that are lost or gained, and the final charge of the ion that is formed. 5

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7 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11.

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9 Names of columns on the Periodic Table alkali metals noble gases alkaline earth metals halogens transition metals What makes the Noble Gases noble If an atom has every orbital in its valence level filled by two electrons, it also carries the least amount of energy possible for it to carry. Atoms form ions and bond to each other in order to come as close as possible to having their valence orbitals filled. Noble Gases don t bond to anything (don t react with anything) because they already have their valence level filled bonding would only increase their energy. Just like kings and queens from the past, the Noble Gases don t have to do anything and everyone wants to be like them. Things Fall Down And Break Memorize this: everything seeks lower energy and greater chaos. It s called the Second Law of Thermodynamics. Everything in the universe tends toward lower energy and greater randomness. All chemical reactions are for the sake of gaining lower energy or greater chaos. The universe will not allow reactions which both increase energy and decrease chaos. Predicting the Charge on Ions Atoms will steal electrons from each other or give electrons to each other in order to have their valence level resemble that of a Noble Gas. For example, fluorine will steal one electron and end up resembling neon. It won t BE neon because neon has 10 protons and fluorine has 9 protons.

10 +1 However, after the theft, fluorine will have 10 electrons just like neon. There will be one other big difference between them (besides the number of protons) neon is neutral while fluorine will have a -1 electrical charge. Once an atom changes its number of electrons and gets a charge, it is called an ion. Negative ions have their names end in ide. So, technically, after fluorine gains an extra electron it would be called fluoride ion. The reason fluoride has a charge is that it has 10 e - but only 9 protons (+). So, one of the negative charges isn t canceled out by a positive charge. total electron charge -10 total proton charge +9 total combined charge -1 Charge computation for fluoride ion. Since all atoms are neutral until they gain or lose electrons, the number of electrons they gain or lose becomes their charge. For instance, oxygen gains two electrons, becoming oxide ion with a minus two charge. Aluminum loses three electrons, becoming aluminum ion with a plus three charge. total electron charge -10 total proton charge +13 total combined charge +3 Charge computation aluminum ion. Atoms become ions in order to resemble a Noble Gas. So, in the example of aluminum above, Al can either lose 3 electrons and resemble Ne or gain 5 electrons and resemble Ar. Atoms will always choose the shorter path, so Al loses 3 electrons. While this sounds a little complicated, ion formation follows some obvious trends where all ions in a column on the Periodic Table have the same charge. skip: too complex for physical science +2 +3-3 -2-1 The Periodic Table, ideally, would be a Periodic Spiral since Ne should touch Na, Ar should touch K, etc. doesn t form ions

11 Ions Over the Aeons Complete the table below. One of these will be 0. atomic # symbol # of valence electrons # of e - lost in ion formation # of e - gained in ion formation charge on ion 1 3 4 5 7 8 9 11 12 13 14 15 16 17 53 56

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13 Ions Ions Ions Choose the best answer. 1) What is the charge on neutral hydrogen? a) -1 b) 0 c) 1 d) 2 2) What is the charge on an Mg ion? a) -1 b) 0 c) 1 d) 2 3) What is the charge on F ion? a) -1 b) 0 c) 1 d) 2 4) What is the charge on H ion? a) -1 b) 0 c) 1 d) 2 5) What is the charge on calcium ion? a) -1 b) 0 c) 1 d) 2 6) What is the charge on I ion? a) -1 b) 0 c) 1 d) 2 8) What is the charge on Na ion? a) -1 b) 0 c) 1 d) 2 9) What is the charge on O ion? a) -2 b) -1 c) 2 d) 3 10) What is the charge on Al ion? a) -2 b) -1 c) 2 d) 3 11) What is the charge on Br ion? a) -2 b) -1 c) 2 d) 3 12) What is the charge on S ion? a) -2 b) -1 c) 2 d) 3 13) What is the charge on Be ion? a) -2 b) -1 c) 2 d) 3 7) What is the charge on neutral N? a) -1 b) 0 c) 1 d) 2

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15 Lewis Diagrams of Ions When drawing Lewis diagrams of ions, always put brackets around the symbol for the element and then on the outside right of the brackets include a superscript of the ion s charge. For example, Al 3+ would be drawn as [Al] 3+. Often the valence electrons of the final ion are included:.. 3+ [ ]... Al... Remember, aluminum had 3 valence electrons. It lost them. Now its valence structure looks like that of the Noble Gas, neon (which was the whole point of the ionization process).

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19 Electronegativity Electronegativity is a measure of how strongly an atom pulls a shared electron pair towards it. The greater the electronegativity of an element, the stronger its attractive pull on electrons. For example, in a molecule of hydrogen bromide (HBr), the electronegativity of bromine (2.8) is higher than that of hydrogen (2.1), and so the shared electrons will spend more of their time closer to the bromine atom. Bromine will have a slightly negative charge, and hydrogen will have a slightly positive charge. In a molecule like hydrogen (H2) where the electronegativities of the atoms in the molecule are the same, both atoms have a neutral charge. In general, electronegativity can be predicted: F is the most electronegative element and the closer an element is to F, the more electronegative it is. O and Cl are more electronegative than is S. The biggest difference in electronegativity is The concept of electronegativity was introduced by Linus Pauling in 1932, and this became very useful in predicting the nature of bonds between atoms in molecules. In 1939, he published a book called The Nature of the Chemical Bond, which became one of the most influential chemistry books ever published. For this work, Pauling was awarded the Nobel Prize in Chemistry in 1954. He also received the Nobel Peace Prize in 1962 for his campaign against above- ground nuclear testing. between metals and non-metals. Electronegativity is important because the difference in electronegativity of two atoms bonded to each other determines the type of chemical bond. It (along with the geometry of the final molecule) also determines whether or Polar and non-polar bonds If two atoms are bonded to each other and the two atoms have a large difference in electronegativity, the one with the greater electronegativity will become negatively charged and the less electronegative atom will become positively charged. When this happens, the bond they form is said to be polar. If the atoms have nearly equal electronegativities, no charges form around the atoms and the bond is said to be non-polar Polar and non-polar molecules An atom with exposed positive and negative regions is said to be polar An atom with no exposed positive or negative regions is said to be non-polar. All molecules with polar bonds are not necessarily polar molecules. Geometry also plays a role and can hide the positive and negative regions inside the shape of the molecule so that the molecule effectively is non-polar. Polar molecules always have polar bonds. Polar bonds in a molecule do not always result in polar molecules.

20 not the molecule formed by those atoms has charged regions (which are very important for chemistry). Bonding In the model of the atom that has been used so far, the atom is made up of a central nucleus, surrounded by electrons that are arranged in energy levels (also called shells). The shape of the paths the electrons take are called orbitals. Each energy level is made of a group of orbitals, all extending roughly the same distance from the nucleus. The electrons in the outermost energy level of an atom are called the valence electrons. This model of the atom is useful in trying to understand how different types of bonding take place between atoms. You will remember from these earlier discussions of electrons and energy levels in the atom, that electrons always try to occupy the lowest possible energy level. In the same way, an atom also prefers to exist at the lowest possible energy state so that it is most stable. An atom is most stable when all its valence electron orbitals are full. In other words, the outer energy level of the atom contains the maximum number of electrons that it can. A stable atom is also an unreactive one, and is unlikely to bond with other atoms. This explains why the noble gases are unreactive and why they exist as atoms, rather than as molecules. Look for example at the electron configuration of neon (1s 2 2s 2 p 6 ). Neon has eight valence electrons in its valence energy shell. This is the maximum that it can hold and so neon is very stable and unreactive, and will not form new bonds. Other atoms, whose valence energy levels are not full, are more likely to bond in order to become more stable. We are going to look a bit more closely at some of the energy changes that take place when atoms bond. Energy and bonding Let s start by imagining that there are two hydrogen atoms approaching one another. As they move closer together, there are three forces that act on the atoms at the same time. These forces are shown below:

In the example of the two atoms, when the atoms are closer together the atoms are pulled towards each other. As this happens, the potential energy of the system decreases (because energy would now need to be supplied to the system in order to move the atoms apart). However, as the atoms move closer together, repulsive forces start to dominate and this causes the potential energy of the system to rise again. At some point, the attractive and repulsive effects are balanced, and the energy of the system is at its minimum (point X). It is at this point, when the energy is at a minimum, that bonding takes 21

22 place. The distance marked P is the bond length, i.e. the distance between the nuclei of the atoms when they bond. Q represents the bond energy (bond strength) i.e. the amount of energy that must be added to the system to break the bonds that have formed. Bond strength means how strongly one atom attracts and is held to another. The strength of a bond is related to the bond length, the size of the bonded atoms and the number of bonds between the atoms. In general, the shorter the bond length, the stronger the bond between the atoms, and the smaller the atoms involved, the stronger the bond. The greater the number of bonds between atoms, the greater will be the bond strength. What happens when atoms bond? A chemical bond is formed when atoms are held together by attractive forces. This attraction occurs when electrons are shared between atoms, or when electrons are exchanged between the atoms that are involved in the bond. The sharing or exchange of electrons takes place so that the outer energy levels of the atoms involved are filled and the atoms are more stable. If an electron is shared, it means that it will spend its time moving in the electron orbitals around both atoms. If an electron is exchanged it means that it is transferred from one atom to another, in other words one atom gains an electron while the other loses an electron. The type of bond that is formed depends on the elements that are involved. In this class, we will be looking at three types of chemical bonding: covalent, ionic and metallic bonding. You need to remember that it is the valence electrons that are involved in bonding and that atoms will try to fill their outer energy levels so that they are more stable. Ionic Bonding The nature of the ionic bond Ionic bonding takes place when the difference in electronegativity between the two atoms is more than 1.7 (a chart is used to determine electronegativities). This usually happens when a metal atom bonds with a non-metal atom. When the difference in electronegativity is large, one atom will attract the shared electron pair much more strongly than the other atom does, eventually stealing or nearly stealing the electron. This transfer of an electron also causes a transfer of negative charge since the electron has a -1 charge. So, the atom which gained the electron becomes negative. The atom which lost the electron becomes positive. Positive and negative attract, so the atoms (now ions because they have a charge) stick to each other. Example 1: In the case of NaCl, the difference in electronegativity is 2.1. Sodium has only one valence electron, while chlorine has seven. Because the electronegativity

23 of chlorine is higher than the electronegativity of sodium, chlorine will attract the valence electron in the sodium atom very strongly. This electron from sodium is transferred to chlorine. Sodium has lost an electron and forms a Na + ion. Chlorine gains an electron and forms a Cl ion. The attractive force between the positive and negative ion is what holds the molecule together. The balanced equation for the reaction is: Na + Cl NaCl This can also be represented using x s for chlorine s valence electrons and dots for sodium s valence electrons.: Example 2: Another example of ionic bonding takes place between magnesium (Mg) and oxygen (O) to form magnesium oxide (MgO). Magnesium has two valence electrons and an electronegativity of 1.2, while oxygen has six valence electrons and an electronegativity of 3.5. Since oxygen has a higher electronegativity, it attracts the two valence electrons from the magnesium atom and these electrons are transferred from the magnesium atom to the oxygen atom. Magnesium loses two electrons to form Mg 2+, and oxygen gains two electrons to form O 2. The attractive force between the oppositely charged ions is what holds the molecule together. The balanced equation for the reaction is: 2Mg + O2 2MgO Because oxygen is a diatomic molecule (always comes in pairs), two magnesium atoms will be needed to combine with two oxygen atoms to produce two molecules of magnesium oxide (MgO). Three or more atoms can also bond together. For example, Aluminum forms a 3 + ion and chlorine forms a 1 - ion. Since the ions will combine in a way which cancels out the charge, three chlorines will combine with one aluminum to make AlCl3. [Cl] - [Cl] - [Al] 3+ [Cl] -

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27 Lewis Diagrams of Ionic Compounds Since ionic compounds are merely positive and negative ions stuck together, the Lewis Diagram is merely the diagram of the ions beside each other. Unless your instructor says otherwise, for ionic compounds you do not need to show the dots for the electrons. Aluminum and chlorine combine: Sodium and oxygen combine: [Cl] - [Cl] - [Al] 3+ [Cl] - [Na] + [O] 2- [Na] + It doesn t matter exactly which position the ions have. All that matters is that the ions are bonded on the top, bottom, left, or right and that same charges don t bond with each other. For example, sodium and oxygen could be written this way: [Na] + [O] 2- [Na] + The nature of the metallic bond The structure of a metallic bond is quite different from covalent and ionic bonds. In a metal bond, the valence electrons are delocalized, meaning that an atom s electrons do not stay around that one nucleus. In a metallic bond, the positive atomic nuclei are surrounded by a sea of delocalized electrons which are attracted to the nuclei. Metallic bonding is the electrostatic attraction between the positively charged atomic nuclei of metal atoms and the delocalized electrons in the metal. We will not deal further with these in physical science.

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29 Ionic Bonding Draw the Lewis diagrams of the ionic compounds which result from the combining of the listed elements. Lewis Diagram N, K N, Ca Ca, I Al, O F, Sr Cs, F Sr, S Rb, S

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31 Ionic Bonding Model Lab Use the ball and stick model set your group is given to model the following compounds. The elements are provided below. Write the correct formula, draw the ionic Lewis Diagram and then model them. Show all unbonded electrons as sticks. Model positive ions valence electrons from their final (ionized) structure. When you are done with the first seven, show them to your instructor. Then complete the rest. 1) Na, Cl 2) K, O 3) Mg, O 4) Li, N 5) Ba, I 6) Al, Br 7) Fe 2+, O 8) Pb 4+, I 9) Pb 4+, O

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33 periodic trends review Name: Period: 1. Refer to the periodic table. How many protons does the element chlorine have? A. 7 B. 17 C. 18 D. 35 2. How many valence electrons do alkaline earth metals have? A. 0 B. 1 C. 2 D. 3 3. An atom of calcium has how many valence electrons? A. 2 B. 3 C. 20 D. 40 4. Which of the following is a metalloid? A. sodium B. copper C. silicon D. iron

34 5. What does the very left group of the periodic table contain? A. The most reactive metal B. The least reactive metals C. The most reactive nonmetals D. The least reactive nonmetals 6. What kind of chemical bond is formed when a transfer of electrons occurs? A. ionic B. hydrogen C. covalent D. metallic 7. Choose the compound that is not covalently bonded. A. H 2 B. CO 2 C. H 2 O D. NaCl 8. Chemical bonds form in order for the chemicals involved to have energy. A. lower B. higher 9. Which bond type could best be called the sea of electrons? A. covalent B. metallic C. ionic

Covalent Bonding Covalent bonding occurs between atoms of non-metals. The outermost orbitals of the atoms overlap so that unpaired electrons in each of the bonding atoms can be shared. By overlapping orbitals, the outer energy shells of all the bonding atoms are filled. The shared electrons move in the orbitals around both atoms. As they move, there is an attraction between these negatively charged electrons and the positively charged nuclei, and this force holds the atoms together in a covalent bond. Covalent bonding is a form of chemical bonding where pairs of electrons are shared between atoms. 35

The above examples all show single covalent bonds, where only one pair of electrons is shared between the same two atoms. If two pairs of electrons are shared between the same two atoms, this is called a double bond. A triple bond is formed if three pairs of electrons are shared. 36

Lewis notation and molecular structure Although we have used diagrams to show the structure of molecules, there are other forms of notation that can be used, such as Lewis notation. Lewis notation uses dots and crosses to represent the valence electrons on different atoms. The chemical symbol of the element is used to represent the nucleus and the core electrons of the atom. 37

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45 bonding name: period: 1. What kind of chemical bond is formed when a transfer of electrons occurs? A. ionic B. hydrogen C. covalent D. metallic 2. How is a covalent bond formed? A. bond between a metal and a non-metal B. transfer of electrons between atoms C. the sharing of electrons between two atoms D. attraction between a positive metal ion and a shared sea of electrons 3. What does the Lewis dot formula for Br show? A. a single ionic bond B. a single covalent bond C. a double covalent bond D. a triple covalent bond

46 4. What are the kinds of bonds found on carbon in CO - ²? A. 3 single bonds B. 3 double bonds C. 1 single and 2 double bonds D. 2 single and 1 double bonds