SCH3U- R. H. KING ACADEMY ATOMIC STRUCTURE HANDOUT NAME:

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1 Particle Theory of Matter Matter is anything that has and takes up. All matter is made up of very small. Each pure substance has its of particle, from the particles of other pure substances. Particles each other. Particles are always. Particles of a higher temperature move, on average than particles at a temperature. Dalton s theory consisted of the following five postulates. All matter is composed of tiny indivisible particles called. All atoms of an element have properties. Atoms of different elements have properties. Atoms of two or more elements can combine in definite number ratios to form. Atoms cannot be nor. LAW OF CONSERVATION OF MASS The mass of the reactants in a chemical reaction is to the mass of the products. This is known as the LAW OF CONSERVATION OF MASS. Atoms are just rearranged into different compounds. Using this idea, solve the following problems 2KClO 3 2 KCl + 3O 2 If 500 g of KClO 3 decomposes and produces 303 g of KCl, how many grams of O 2 are produced? Answer: LAW OF CONSTANT COMPOSITION Dalton also proposed that atoms of different elements can be chemically combined in a number ratio to form compounds. This is known as the LAW OF CONSTANT COMPOSITION. For example: A compound is a substance that has a, composition and properties different from the that form it. Molecules of a compound will always contain the elements present in the ratio. If the elements combine in a different ratio then a compound is formed. H1

2 HISTORICAL OUTLINE - MODEL OF THE ATOM It has been determined that the atom contains three basic sub-atomic particles. For each determine the following information. Particle Date of Name of the Brief description of the experiments that led to discovery. Scientist credited their discovery. with the discovery. Electron Proton Neutron Elements are represented by the symbols found on the periodic table. The of the element is written as a subscript before the symbol. The is written as a superscript before the symbol Example: Atomic number & Atomic Mass The atomic number of the element represents the number of in the element and is equal to the number of in a atom. Example: The atomic mass number of the element represents the number of plus the number of in the atom. The number of neutrons in an element can be determined by subtracting from atomic. H2

3 ATOMS Each element is made up unique atoms. The atoms of an element have identical properties All neutral atoms of an element may have the SAME NUMBER of AND Some atoms of an element may have a DIFFERENT NUMBER of and therefore different. ISOTOPES Atoms of the same element having the same number of protons in their nucleus but a different number of neutrons are of each other. The average atomic mass number shown in the periodic table is the result of an average mass based on the of each isotope. Most elements occur naturally as mixtures of isotopes, as indicated on your handout. The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. AVERAGE ATOMIC MASS Isotopes with different mass numbers exist in a fixed ratio in a sample of an element. The percent abundance of each isotope can be determined by mass. This percent abundance is used to calculate the average atomic mass of the element. It is used as a weighted measure of the mass of a specific isotope. AVERAGE ATOMIC MASS [A and B are isotope masses] AAM = % abundance x mass A + % abundance x mass B Example: A sample of carbon has two isotopes C-12 and C-13, with C-12 comprising % of the sample and C-13 comprising 1.11 %. Find the average atomic mass. [Ans: u] Question: Sometimes an isotope is written without its atomic number - e.g. 35 S (or S-35). Why? Example: A sample of two naturally occurring isotopes of lithium, Li-6 and Li-7 have masses of 6 u and 7 u, respectively. Which of these two occurs in greater abundance? H3

4 Some isotopes are stable like H-2 while other isotopes are unstable like C-14. Unstable isotopes are called. Radioisotopes undergo radioactive resulting in the production of ionizing radiation and a more stable nucleus. Each radioisotope has a characteristic rate of decay that is known as a. HALF - LIFE Radioisotopes undergo radioactive decay at unique rates that are characteristic for each different radioisotope. The time it takes for of the nuclei in a radioactive sample to decay is known as the half-life of the radioisotope. Half-lives may vary from a few seconds to many years. [Po-226 has a half-life while Cs-142 has a a (years) half-life] CARBON DATING Uses to date organic material. Carbon- 14 has a half-life of. Carbon is constantly through the carbon cycle through living organisms, and the proportion of carbon- 14 remains. Once an organism dies, the recycling of carbon and C -14 starts to. As time goes on the remains contain fewer and fewer C -14 atoms. By the amount of C -14 in an organic sample to the amount present in organisms, it is possible to determine the age of the organic sample. Uses of radioisotopes Non living material such as rocks can be dated similarly using with a half-life of 1.3 x 10 9 a. Anthropologists and geologists commonly use these techniques to date both once as well as rocks. Other disciplines that find these techniques useful are,, and. Half - Life Problems Example: Phosphorus-32 has a half-life of 14.3 days. How many mg of phosphorus-32 remain after 57.2 days if you start with 4.0 mg of the isotope? H4

5 Another way to work out half- life problems: 1 t t 1 2 Another way to work out half-life problems: Use a Formula Af = A i 2 A f is the amount of substance left A i is the initial amount of substance t is the elapsed time t 1/2 is the half-life of the substance RADIOISOTOPES Unstable isotopes undergo radioactive giving off radiation and changing the composition of their. This emission of radiation from the nucleus of an atom is known as. There are 3 types of radiation given off by radioisotopes. 1. particles ( 4 2He 2+, ) - the of helium atom - have protons and neutrons - have a charge - stopped n a few centimeters of 2. particles (, e - ) - electrons travelling at a very high speed - have a negative charge - stopped by a thin sheet of or foil. 3. rays ( ) -do not consist of - a form of high energy radiation, to X-rays - stopped by a sheet of several centimeters thick or reinforced H5

6 The Modern Periodic Table 1. An arrangement of the elements in order of their numbers so that elements with properties fall in the same column (or group). Groups: vertical columns (#1-18) Periods: horizontal rows (# 1-7) 2. Periodicity the of the elements in the same group is explained by the arrangement of the around the nucleus. Chemical Families: 1. Group 1: Alkali metals silvery metals; most of all metals, never found free in nature; reacts with to form alkaline or basic solutions; stored under ; whenever you mix Li, Na, K, Rb, Cs, or Fr with water it will and produce an alkaline solution 2. Group 2: Alkaline earth metals reactive than Alkali, but still react in water to produce an solution; never found in nature; harder, denser, stronger than alkali; 3. Transition Metals They are all with properties (malleability, luster, good conductors, etc ); are referred to as the and than alkali or alkaline; less than alkali or alkaline; (Ex: Ni, Pd, Pt) Groups Contain and ;, along zigzag line, have characteristics of metals and nonmetals (many are conductors but are ). The metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. Group 17 - Halogens most reactive nonmetals electrons in outermost energy level (that is why so only need one electron to have 8); called the formers (they react vigorously with metals to form salts). A salt is a ion and a ion bonded together. Most are. Group 18 - Noble gases- unreactive All are. Inner Transition Metals Lanthanides shiny metals; Ce-Lu Actinides and ; Th-Lr Hydrogen and Helium - Oddballs Hydrogen is NOT an Alkali metal; it is a very gas. It is placed with the Alkali metals because it only has 1 valance electron. Helium is a Noble gas, it is, but it does not have 8 electrons in outermost energy level, because it only has 2 total electrons! H6

7 Trend in Atomic Radii Atomic Radius Atomic radius is the distance from the centre of the nucleus of an atom to the outermost electron. The size of an atomic radius cannot be measured exactly because it does not have a defined boundary. However the atomic radius can be thought of as the distance between the nuclei of identical atoms joined in a molecule. The greater the number of energy levels the is the distance of the outermost electron to the center of its atom s nucleus. Group trend - atomic radii as you move up a group. Period trend atomic radii as you move across a period. Ionic Radius Ionic radius is the distance from the centre of the nucleus of an to the outermost electron. Cations will have a ionic radius than the neutral atom. Anions will have a ionic radius than the neutral atom. Force of Attraction: The force of attraction between negatively charged electrons and the positively charged nucleus is the attraction of opposite charges. The force of attraction existing between the outermost electron and the middle of the nucleus is dependent on two factors: 1. The of the positive charge - determined by the number of protons in the nucleus. 2. The between the outermost electron and the nucleus. A balance exists between the of the electrons to the nucleus and the of the electrons between themselves Ionization Energy (IE) Ionization Energy is the energy in kilojoules per mole (kj/mol) needed to the outermost electron from a gaseous atom to form a ion (cation). Na + Energy Na + + e - NOTE: Metals react to electrons. The stronger an electron is held the the IE needed to ionize (pull away) that electron Trend: Ionization Energy (IE) Group trend ionization energy as you move a group (or as you move a group). Period trend Ionization energy as you move across the period (left to right). Example: Which atom has the higher first ionization energy? Hf or Pt. Explain. Example: Which atom has the higher first ionization energy? Cl or Ar. Explain. H7

8 Electron Affinity [EA] Electron affinity is the energy in kilojoules per mole (kj/mol) when an electron is by an atom to form a ion (anion). Cl + Electron Cl - + Energy NOTE: Nonmetals react to electrons Trend: Electron Affinity (EA) Period trend electron affinity as you move across a period because atoms become and the nuclear charge. This means there is a greater pull from the nucleus. Group trend electron affinity as you move a group (or as you move a group) because the size of the atom. Example: Which element has the greater electron affinity? Pb or Sn. Explain Electron affinity vs. Ionization energy Electron affinity and Ionization energy follow the same trend in the periodic table. The the attraction an atom has for electrons the it will be to remove electrons from that atom and the the IE energy will be. The the attraction for electrons the the energy released when an atom gains an electron. Electronegativity [EN} Electronegativity is a measure of the of an atom to gain electrons when it is chemically combined (bonded) to another element. The the pull or attraction of electrons to an atoms nucleus, the its tendency to gain electrons. In general, metals have EN and nonmetals have EN. The actual amount of EN an atom has is indicated by a number of the Pauling Electronegativity Scale that goes from 0 to 4. Dr. Linus Pauling set up this scale and gave the element having the greatest EN an arbitrary number of 4, and he assigned numbers to the others relative to this element. Trend: Period trend - EN as you go across a period (excluding the noble gases) because size. Group trend - EN as you go a group because there is pull from the nucleus as the electrons get further away. Example: Which would have the greater EN? Ca or Se. Explain. Electronegativity enables us to predict what of bond will be formed when two H8

9 elements combine. Reactivity of Elements Reactivity - how a substance reacts with another. Metals lose electrons (Ionization Energy) Nonmetals gain electrons (Electron Affinity) Trends for Metals: Group Trend: Period Trend: Trends for Nonmetals: Group Trend: Period Trend: EFFECTIVE NUCLEAR CHARGE (ENC) AND SHIELDING Element Na Mg Al Si P S Cl Ar # of electrons # of valance electrons # of protons # of inner electrons ENC EFFECTIVE NUCLEAR CHARGE (ENC) AND SHIELDING The force of attraction between charged protons in the nucleus and charged electrons is the force that holds atoms together. The inner electrons (not in the outermost energy level) in inner energy levels, partially or the attraction of the protons from the outer electrons in the outermost energy level (VALENCE ELECTRONS). H9

10 The canceling of the positive nuclear charge is called. EFFECTIVE NUCLEAR CHARGE (ENC) is a number assigned to elements to describe the amount of shielding by the valence electrons. ENC = Number of - Number of electrons The greater the ENC the the valence electrons are shielded and the the pull on the valence electrons. Greater ENC will mean a atomic radius. Shielding will help explain some of the trends in the periodic table SUCCESSIVE IONIZATION ENERGIES The first ionization energy is the energy required to remove the electron (First IE). It is relatively low because of the repulsion exerted by the other electrons. Each successive Ionization energy (Second and Third IE and so on) will. It becomes difficult to remove successive electrons since the pull of the nucleus becomes (greater number of protons relative to the electrons) and the electrons are tightly held Ionic radius becomes. There will be a jump in the increase of IE once the gas configuration has been reached. This is because outer energy level has been (radius is smaller) Example 1: Consider the following Ionization Energies for an element X: How many valance electrons does this element have? Explain. 1 st 2 nd 3 rd 4 th 5 th 2.38 kj 2.54 kj kj kj kj Example 2: Where would the large increase in I.E. occur for Se? Explain your answer. H10

11 PROPERTIES OF ATOMS The elements in the Periodic Table are arranged according to atomic number into groups and rows. Elements in a group have similar chemical properties. If the element is a member of a "long" group, then the group number indicates the number of electrons in the outer energy level of that element. The row number indicates the number of occupied energy levels that are found about the nucleus of the atom. Bohr-Rutherford diagrams are particularly useful when discussing the trend in atomic properties in the Periodic Table or comparing two elements and their respective values for an atomic property. Force of Attraction The force of attraction that exists between an outermost electron and the nucleus of the atom is dependent upon two factors as illustrated in Coulomb's Law. These two factors are a} the size of the positive charge that the electron experiences b} the distance between the outermost electron and the nucleus. A) The positive charge in an atom is determined by the number of protons in the nucleus of the atom. The outermost (valence) electrons being negatively charged experience an attraction by this positively charged nucleus. At the same time the electrons in the inner energy levels repel the valence electrons since both are negatively charged. The net result between the attraction of the nucleus and the repulsion by the inner electrons is called the effective nuclear charge (ENC). Very "roughly" the ENC can be said to be equal to the [actual nuclear charge (number of protons in the nucleus)] minus [the number of electrons in the inner energy levels. For example for oxygen the ENC = +6 where the ENC = +2 for calcium. Draw B-R diagrams of both elements to verify these values. Thus as the ENC increases so does the force of attraction between the valence electrons and the nucleus. This being a direct relationship as one increases so does the other and vis versa. B) The distance between the valence electrons and the nucleus can be said to be directly related to the number of energy levels occupied around the nucleus. In this case the force of attraction is inversely related to the square of the distance between the two. Simply the further away the electron is from the nucleus the weaker the force of attraction. Likewise the closer to the nucleus (shorter distance) the greater the attraction. In summary: (i) the larger the charge experienced by the particles the greater the force of attraction (ii) the closer the particles are to each other, the stronger the force of attraction. PERIODICITY Periodicity, the concept, that the trend of these properties was repeating, was noted back in the nineteenth century and led to the establishment of the periodic table in the form that we know it today. IONIZATION ENERGY: The energy required to remove the outermost electron from an atom to form an ION. An ion is an atom having either an excess of electrons (negative ion) or a shortage of electrons (positive ion). The stronger the electron is held the (greater force of attraction) the larger the energy required to remove an electron. H11

12 After the outermost electron in any atom is removed (First Ionization Energy) successive ionization energies (Second I.E., Third I.E. etc.) increase as it becomes more difficult to remove electrons which are closer to the nucleus (stronger attraction). ELECTRON AFFINITY: The energy released when an atom gains an electron to form a negative ion. Electron affinity is the opposite of Ionization Energy. The stronger the affinity (attraction) that an atom has for it's electrons, the harder it will be to ionize (pull away) that electron. The higher the ionization energy, the greater the electron affinity (direct relationship). The greater the affinity the larger the amount of energy released when an atom gains an electron. ELECTRONEGATIVITY: A measure of an atom's tendency to gain electrons. An atom which is highly electronegative means that it has a very strong pull (attraction) on it's electrons. Electronegativity and Ionization Energy are also directly related. ATOMIC RADIUS: The distance from the centre of the nucleus to the outermost electron. The units of measurement are picometres (10-12 m) eg. Nitrogen radius: 75 pm. Therefore for each atom, the more energy levels filled by electrons, the greater the atomic radius. Explanations of the Trends: For example: Going down a particular group: Because of the increased number of energy levels, atomic radius increases. The ENC remains constant. Therefore, force of attraction (F of A) between the nucleus and the outer electrons should decrease as one goes down the group. Therefore it will be easier to lose (remove) an electron and there will be a lower tendency to gain an electron. Therefore the IE will decrease and the EN will decrease as one goes down the group. The Atomic Radius will increase due to the increased number of occupied energy levels as one goes down the group. Going across a particular row from left to right: Although each atom has the same number of energy levels, the radius decreases due to the increased nuclear charge "pulling the electrons in". The ENC increases. Therefore force of attraction (F of A) between the nucleus and the outer energy level electrons should increase. Therefore it will become more difficult to remove an electron and the tendency to gain electrons will become stronger. Therefore IE will increase and the EN will increase as one goes across a row of the periodic table from left to right. Example: Explain why Al has a higher IE and EN than Mg. Al (13p, 14n) 2e 8e 3e ENC = +3 Mg (12p, 14n) 2e 8e 2e ENC = +2 Although Al and Mg have the same number of energy levels, the radius of Al will be smaller than the radius of Mg due to the increased nuclear charge. Force of attraction (F of A ) in (Al) is greater than F of A in (Mg) due to smaller radius and larger enc. Therefore it is more difficult to remove an electron from Al. Also there is a stronger tendency to gain electrons in Al. Therefore IE (Al) > IE (Mg) and EN (Al) > EN (Mg) H12