Chapter 3: Chemical Foundations Elements, Atoms, and Ions 3-1
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. 3-2
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. 3-3
Objectives Predict which ion a given element forms by using the periodic table. Describe how ions combine to form neutral compounds. 3-4
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. 3-5
Table 3.1 3-6
are made up, mainly, of oxygen, carbon, hydrogen and nitrogen. 3-7
Section 3.2: Symbols For The Elements Just as each state has a two-letter abbreviation, each element has a one- or two-letter symbol to make life simple for chemists. 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. 3-8
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). 3-9
are well over 100 different elements, many are fairly rare; we should know the most common elements. 3-10
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. 3-11
Dalton s Atomic Theory 1. Elements are made of tiny particles called atoms. 2. All atoms of a given element are identical. 3. The atoms of a given element are different from those of any other element. 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms. 5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way atoms are grouped together. 3-12
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) C 6 H 12 O 6 or H 3 PO 4 3-13
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. 3-14
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). 3-15
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. 3-16
Figure 3.3: Plum Pudding model of an atom. 3-17
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. 3-18
Figure 3.5: Rutherford s experiment. 3-19
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. 3-21
Figure 3.6: Actual results. 3-22
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. 3-24
Figure 3.9: A nuclear atom viewed in cross section. 3-25
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. 3-26
. Section 3.7: Isotopes isotopes; two atoms of the same element (same number of protons) with different numbers of neutrons. 3-27
Figure 3.10: Two isotopes of sodium. 3-28
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. 3-29
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: A X Z 3-30
Isotopes The two previous examples of isotopes of sodium would be: 23 Na 24 Na 11 11 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). 3-31
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. 3-32
Section 3.8: Introduction to the Objectives: Periodic Table 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 3-34
In any box on the Periodic Table, what information can you find? 6 C 12.01 Average Atomic Mass = the weighted average of all the mass numbers for each isotope of the element 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 3-35
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 3-36
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 3-37
Practice Oxygen has 3 isotopes 16 O, 17 O, 18 O 99.76% of mass is 16 O 0.04% of mass is 17 O 0.20% of mass is 18 O What is the average atomic mass? Find the atomic mass if 99.64% of mass is 14 N and 0.36% is 15 N. 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? 3-38
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 3-39
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 3-40
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. 3-41
Figure 3.12: Elements classified as metals and nonmetals. 3-42
Metals: Fall to left and under the stairs Properties of Metals: Efficient conduction of heat and electricity Malleability Ductility A lustrous appearance Positively charged ions 3-43
Non-Metals: Right and above the stairs Dull, Brittle Negatively charged ions Nonconductors -insulators 3-44
Semimetals or Metalloids: Makeup the stairs Properties of both metals and nonmetals Semiconductors 3-45
Lanthanide and Actinide Series Mostly human made elements Radioactive elements 3-46
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 3-47
Group 1A Alkali Metals All members react with water to produce same ratio of products: Li + HOH -> LiOH + H 2(g) Na + HOH -> NaOH + H 2(g) K + HOH -> KOH + H 2(g) Rb + HOH -> RbOH + H 2(g) Cs + HOH -> CsOH + H 2(g) As you move down the group: Reactivity with water increases Melting points and boiling points decrease Density increases 3-48
Group 2A Alkaline Earth Metals Members of Group 2A Form +2 charges as ions React with water 3-49
Group 2A Alkaline Earth Metals The members of this group also react to form the same ratio: Mg + HOH -> Mg(OH) 2 + H 2 Ca + HOH -> Ca(OH) 2 + H 2 Sr + HOH -> Sr(OH) 2 + H 2 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 3-50
+1 +2 Trends for Group 1A & 2A Melting/Boiling points increase Reactivity increases Density increases 3-51
Group 3A or 13 Form +3 charge as ions React with water in a 1metal:3 OH ratio Slow reaction with water 3-52
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 3-53
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 3-54
Group 6A or 16: The Oxygen Family Form -2 charge as ions Members below oxygen form oxides with putrid odor: SO 2, SeO 2, & TeO 2 3-55
Group 7A or 17: Halogen Family Exist as diatomic molecules Form ions with a -1 charge Non-metals 3-56
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: H 2, F 2, Cl 2, Br 2, I 2 Members of this family are toxic and poisonous (exception is H 2 ) 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. F 2 + 2Na -> 2NaF explodes Cl 2 + 2Na -> 2NaCl vigorous Br 2 + 2Na -> 2NaBr slower I 2 + 2Na -> 2NaI have to heat to get reaction going So what pattern is in the reactivity as you move down the group? 3-58
Group 7A or 17: Halogens or Hydrogen Family As you move down the group: Boiling/Melting point increases Reactivity decreases Density increases 3-59
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. 3-60
Group 8A or 18: Noble Gases As you move down the period: Reactivity decreases Melting point increases Boiling point increases 3-61
Transition Metals: Group 3B-12B Have many electrons that they are able to share, this allows them to have many different ions: Cr 2+, Cr 3+, Cr 5+, Cr 6+ Always lose electrons to make positive ions These varying charges gives variety of colors to same element Malleable and ductile Conduct electricity and Shiny 3-62
+1 Charges of all the Families Remember atoms that lose or gain electrons form ions, and these are the charges each family forms. 0 +2 +3 +4-3 -2-1 Transition Metals 3-63
Figure 3.19: The ions formed by selected members of groups 1, 2, 3, 6, and 7. 3-64
A Few More General Periodic Trends 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? 3-65
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. 3-67
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. 3-68
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 monomeans one. And monatomic gases exist as individual atoms. Figure 3.13: A collection of argon atoms. 3-69
Gases There is another group of gases called diatomic gases. The prefix di- means two. These elements travel in pairs as molecules. Figure 3.14: Nitrogen gas contains NXN molecules. Figure 3.14: Oxygen gas contains OXO molecules. 3-70
Diatomic molecules 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. 3-71
Objectives: 3.10 Ions 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. 3-73
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 - 3-75
Basic Names for Ions Cations do not change names from their neutral atoms. Example: Magnesium loses 2 electrons and becomes Mg 2+ 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. 3-76
Some Common Anion Names What would the names of the following ions be? Chlorine = Fluorine = Bromine = Iodine = Oxygen = Sulfur = 3-77
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. 3-78
Figure 3.19: The ions formed by selected members of groups 1, 2, 3, 6, and 7. 3-79
Practice Determine the name and charge of the following ions: Potassium Bromine Calcium Sulfur Aluminum Strontium Cesium 3-80
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. 3-82
If we add two ions together and they are not equal, then we must add another ion to balance the charge. Example: Mg 2+ + 2Cl - -> MgCl 2 3-83
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? 3-84