Periodic Law Test Questions – Flashcards
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| Triads |
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| Groups of three elements that have similar properties and related atomic masses. |
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| Law of Octaves |
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| An arrangement of elements where every eighth element had similar properties of the first. |
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| Periodic Law |
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| States that the physical and chemical properties of the elements are a periodic funtion of their atomic numbers. |
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| Element |
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| Substances composed of only one kind of atom and cannot be further decomposed through ordinary chemical means. |
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| Atom |
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| The smallest unit of an element that can enter into a combination with other elements. |
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| Diatomic Element |
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| Elements which contain two atoms per molecule. |
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| Group/Family |
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| Element arranged in vertical columns on the periodic chart. |
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| Period/Series |
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| Elements arranged in horizontal rows |
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| Atomic Radius |
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| The distance from the nucleus of the atom to the outermost energy level. |
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| Ion |
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| Atom or group of atoms which possess an overall electrical charge, formed by gaining or losing electrons. |
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| Ionization Energy |
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| The amount of energy required to remove an electron from an atom. |
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| Electron Affinity |
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| The amount of energy released when an electron is added to a neutral atom |
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| Oxidation Number |
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| The number of electrons an atom will gain, lose, or share to form a more stable configuration. |
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| Octet Rule |
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| States that 8 electrons in the outermost energy level is the most stable configuration an atom can have. |
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| Activity |
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| The ability to form ions. |
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| What was the contribution towards the Periodic Chart from Dobereiner? |
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| Responsible for triads. The middle element of the three always had an atomic mass approx. of the other two. |
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| What was the contribution towards the Periodic Chart from Newlands? |
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| If he arranged the elements in order of their increasing atomic mass, there appeared to be a repetition of similar properties every eight elements. Law of Octaves. |
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| What was the contribution towards the Periodic Chart from Mendeleev? |
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| Expanded on Newland's idea. He felt that the properties reoccured after periods of varying length, rather than Newland's 7. He arranged the elements in order of increasing atomic mass and same properties in columns. Table revealed properties repeated in an orderly way. In order to have all elements with similar properties in the same column, he had to leave blank spots- able to predict properties and atomic masses of unknown elements. This enabled scientists to search for and discover missing elements. (Noble Gasses not yet discovered.) |
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| Mendeleev is known as what? |
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| Father of the periodic table. |
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| What was the contribution towards the Periodic Chart from Meyer? |
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| Using the same means, independently, he was able to construct essentially the same periodic table as Mendeleev. He failed to submit documentation of his work. |
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| What was the contribution towards the Periodic Chart from Henry Moseley? |
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| Performed an x-ray experiment which showed how the number of protons in the nucleus varied progressively from element to element. As a result of Moseley's work, the basis for the periodic table was changed to the atomic number instead of Medeleev's atomic masses. |
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| On what basis did Mendeleev arrange the elements in his periodic table? |
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| In order of increasing atomic mass. |
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| On what basis are elements arranged today? |
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| By order of increasing atomic number |
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| Why are two rows of elements located beneath to main body of the periodic table? |
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| To shorten the periodic table. |
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| List the classes of elements and give the properties and location for each. |
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| Metals- shiny, have luster, reflect heat and light, conduct electricity, ductile, malleable. To the left of the stairstep. Non-metals- poor conductors, brittle, non-malleable, non-ductile, dull. Right of Stairs. Metalloids- properties of both, located on stairs. |
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| What must each block on the periodic chart include? |
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| Each block must include only three items- symbol, atomic number, and atomic mass. |
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| A ___ represents one atom of an element. |
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| symbol |
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| An ___ is the smallest part of an element. |
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| atom |
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| List the diatomic elements. |
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| Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine, Astatine |
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| What is the quantative meaning of a chemical symbol? |
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| One atom of an element |
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| How is the periodic talbe formed? |
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| The periodic table is formed by placing elements with similar electron configurations in the same column and listing the elements in order of increasing atomic number. |
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| What group or family of elements was missing from Mendeleev's table? |
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| The Noble Gases Family |
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| What does the group number of the group A elements represent? |
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| Number of valence electrons |
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| Why does the knowledge about one element in a group aid in the understanding of the chemistry of other elements in that group? why? |
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| All elements within a group have the same number and arrangement of valence electrons. Since the number and arrangement determine the chemical properties of an element, all elements within a group have similar properties. |
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| The outermost energy level of the atom can contain a maximum of ___ electrons. These electrons are located in the ___ sublevels. |
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| 8. S & P. |
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| What information is given by the period number? |
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| The number of energy levels and the outermost energy level of the element. |
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| What determines the ionization energy of an atom? |
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| It is determined by the atomic attractive force. The stronger the atomic attractive force, the higher the ionization energy. |
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| Explain the trend for ionization energy as related to the periodic chart within groups and within periods. |
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| Group- the number of energy levels increases, increasing the distance between protons and electrons, weater atomic attractice force requiring less energy to remove an electron, ionization energy decreses. Period- more protons and electrons in the same average distance, greater atomic attractive force, requiring more energy to remove an electron, ionization energy increases. |
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| How would you expect the ionization energies of two atoms of about equal size, but different atomic numbers to compare? why? |
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| The ionization energy of the larger atom will be greater because the ionization energy is determined by the atomic attractive force. The stronger the atomic attractive force, the higher the ionization energy. |
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| How are positive ions formed? What class of elements forms positive ions? |
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| Metals for positive ions by losing electrons. |
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| How does the size of a positive ion compare with the size of the atom from which it is formed? Explain. |
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| In positive ions, the number of protons is greater than the number of electrons. This causes the attractice force of the atom to be greater, therefore the radius of its corresponding neutral atom. |
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| Compare second ionization energies to first ionization energies and explain. |
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| They are always higher than first ionization energies. The atomic attractive force within a positive ion is greater than in the neutral atom, requiring more energy to remove a second electron. |
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| How are negative ions formed? What class of elements form negatice ions? |
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| Non-metals form negatice ions by gaining electrons. |
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| How does the size of a negative ion compare with the size of the atom from which it was formed? Explain. |
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| In negative ions, the number of electrons is greater than the number of protons, causing a weaker atomic attractive force than in the neutral atom where the number of electrons and protons is equal. This causes the radius of the negative ion to be larger than the radius of its corresponding neutral atom. |
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| How is electron affinity related to the atomic attractive force? |
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| The greater the attractive force, the more energy the atom has to release to pull in an electron andthe higher the electron affinity. |
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| Explain the trend for electron affinity as related to the periodic chart within groups and within periods. |
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| Groups- increasing number of energy levels, increasing the distance between protons and electrons, weaker atomic attractive force, the atom has less energy it can release to pull in an electrong, decreasing electron affinity. Periods- going across, more protons and electrons in the same average distance, greater atomic attractive force, the atom has more energy it can release to pull in an electron, increasing electron affinity. Exception: Noble Gases. Have an octet and will not release energy to pull in an electron. |
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| Explain the periodic trends within a group for the following charachteristic: Atomic radius, attractive force, ionization energy, electron affinity. |
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| Increasing number of energy levels, increasing radius, increasing distance between protons and electrons weaker atomic attractive force, inonization- less energy required to remove an electron, Ionization energy decreasing. Electron affinity- decreasing the amount of energy an aton can release to pull in an electron. Electron affinity decreases. |
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| Explain the periodic trends within a period for the following characteristics: Attractive force, atomic radius, ionization energy, electron affinity. |
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| Increasing number of protons and electrons in the same average distance, increasing atomic attractive force. Radius contracts--> decrease. *noble gas. Ionization energy- more energy required to remove an electron. Ionization energy increases. Electron afinity- atom has more energy it can release to pull in an electron. Electron affinity iincreases. *noble gas. Noble Gas-Have more protons and electrons in the same average distance, giving them the greatest atomic attractive force within the period. Due to the octet, electrons spread out, making their radius larger and do not need to gain or lose electrons so electron affinity is non-exsistant. |
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| What determines the reactivity of metals? Non-metals? Explain. |
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| Metals- lose electrons. The more active metal will be the one that requires less energy to remove an electron- the atom with the weater atomic attractive force. Non-metals- gain electrons. The more active non-metal will be the one which will gain an electron the easiest. This will be the atom with the greater atomic attractive force- the atom that has more energy can release to pull in an electron. |
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| List the configurations that show stablility from most stable to least stable. |
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| Most stable- octect. Second- a filled sublevel (s2, p6, d10, f14) Third- a half-filled sublevel (s1, p3, d5, f7) |
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| List the properties and location of Hydrogen. |
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| Properties of Halogen and Alkali metals groups. This is why H is considered in a group by itself. Unstable- very reactive. Diatomic. Natural phase is gas. |
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| List the properties and location of the Alkali metals group. |
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| IA Valence electrons= 1, Lose electrons, Attractive force weak, Ionization energy low, Reactivity- reactivity (unstable) |
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| List the properties and location of Alkaline Earth Metals. |
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| IIA Valence electrons= 2, Lose electrons, attractive force weak, Ionization energy low, Reactivity- active (unstable) |
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| List the properties and location of the Halogen family. |
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| VIIA Valence electrons= 7, non-metals-gain electrons, attractive force strong, Electron affinity high, unstable- very reactive, diatomic. |
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| List the properties and location of the Noble Gas family. |
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| VIIIA Group missing from Mendeleev's chart, gases at room temp., characteristic ionization energy- very high, electron affinity- extremely low, octet- very stable. |
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| List the properties and location of Transistion Elements. |
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| Group B 2 valence electrons, group 1B-1 valence electron, "d" sublevel elecrons, called transition elements because the electrons shift around. The electrons change orbitals within sublevels or they change sublevels. |
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| List the properties and location of Rare Earth metals. |
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| Two rows at the bottom of the chart. 2 valence electrons, "f" sublevel electrons, electrons can shift, elements with more than 92 protons are man-made by neutron bombardment. |
