Chapter 3 – Atoms and Elements Study Goals Classification of Matter Elements and Symbols The Periodic Table The Atom Atomic Number and Mass Number Isotopes and Atomic Mass Electron Energy Levels Periodic Trends All things around us that we can see or touch or feel but don't see (e.g. air current) are made up of what chemists call matter − the Matter material that makes up a substance. Before we look into the stuffs that make up matter at the atomic level, let us define certain terms that chemists use to describe matter macroscopically. A. Matter Can Be Of A Pure Form. i. Pure matters can be made of a single type of atoms (e.g. copper metal, made of 100% copper), in this case they are called pure elements; ii. or they can be made of 2 or more different types of atoms that are chemically combined in a definite proportion (e.g. water consists of 2 H atoms and 1 O atom), in this case they are called pure compounds. B. Matter Can Also Be Of A Mixture A mixture consists of 2 or more substances that are physically mixed but not chemically combined. Examples: Air is a mixture of several gases: N2, O2, CO2, and traces of other gases. Cup made of brass. Brass is an alloy metal of copper and zinc (an alloy metal is one that is made of 2 or more different metals that are homogeneously mixed together). Tea is a mixture of water and substances (caffeine, minerals, and others) that dissolve from the tea powder. Ocean water (a mixture of water, salt (NaCl) and other dissolved minerals). In a mixture, the different types of substances that are found in it are not chemically combined; as a result they can be separated by physical processes. Ex: Place ocean water in a cooking pan and heat it until all the water molecules have evaporated; left behind is a residue of dry, white powder consisting of salt (sodium chloride) and other minerals. Piles of salt surround an evaporation pond in the valley near Glendale, Arizona. Gold panning separates the heavier gold nuggets from sand and dirt. Homogeneous vs Heterogeneous Mixtures A mixture can be classified as homogeneous or heterogeneous. A) In a homogeneous mixture (also called a solution), two or more substances are mixed together in the same phase. They are uniformly dispersed throughout each other even down to the molecular level. IV saline water is a homogeneous mixture of water A bronze cauldron. and dissolved electrolytes such as sodium chloride. Bronze is a homogeneous mixture of copper and tin. B) By contrast to a homogeneous mixture, a heterogeneous mixture does not have a uniform composition throughout the sample. The uneven texture of the material can often be detected. Heterogeneous mixtures may appear uniform macroscopically but on closer examination (i.e. under the microscope), their composition is not uniform. Blood may appear homogeneous until examined under the microscope, which reveals red and white blood cells. Granite is a heterogeneous mixture that contains discrete regions of different minerals (feldspar, mica, quartz). These minerals do not mix together at the molecular level. Pure Element, Compound, or Mixture ? 1) Carbon dioxide (CO2) gas we exhale 2) A salad dressing of oil and vinegar 3) Herbal tea 4) Water and sand 5) A chocolate chip cookie 6) Vinegar (consists of acetic acid and water) 7) Ice 8) Oxygen (O2) 9) Mecury (Hg) in a thermometer 10) Air in a scuba tank Elements and Their Symbols Elements are primary substances from which all things are built and can not be broken down into simpler substances that still retain their identity. - currently 116 elements are known. Of these, only about 90 are found in nature, the rest have been created in the laboratory. Hg S Cu Fe Al Naming of Elements Many elements are named after planets, mythological figures, minerals, colors, famous people, etc. Elements Source of Name Uranium The planet Uranus Titanium Titans (mythology) Chlorine Chloros (Gr. for greenish-yellow) At room temperature, chlorine is a pale yellow- Iodine Ioeides (Gr. for violet) green gas. It is denser than air and will settle to Californium California the bottom of a flask. Curium Marie Curie Iodine is a dark-gray/purple-black solid that sublimes at room temperature into a purple-pink gas that has an irritating odor. Marie Curie was a Polish-born French chemist and pioneered in the early field of radiology and a two-time Nobel laureate. Chemical Symbols Chemical symbols are either one-letter or two-letter abbreviations: One-Letter Symbols Two-Letter Symbols C carbon Co cobalt S sulfur Si silicon Carbon N nitrogen Ne neon I iodine Ni nickel Nickel Note that in a two-letter symbol, only the 1st letter is capitalized, and the 2nd letter is in lower case. ● Many names and symbols are derived from their ancient Greek or Latin names: Name Symbol Silver necklace copper (cuprum) Cu gold (aurum) Au Gold bar iron (ferrum) Fe At RT, sodium is soft and can be cut with a knife. Exposed to humid air, lead (plumbum) Pb the silver surface is quickly oxidized to a chalky potassium (kalium) K white sodium oxide silver (argentum) Ag powder. To prevent exposure to moisture, sodium (natrium) Na sodium is stored often under oil. Physical Properties of Elements Every element has physical properties that include shape, color, odor, taste, density, hardness, melting point, and boiling point. COPPER Physical properties Color reddish-orange Odor odorless Melting point 1083C Boiling point 2567C State at 25C solid copper kettle Luster very shiny Conduction of electricity excellent Conduction of heat excellent copper tubes (Main Group Elements) Periodic Table of Elements Elements Periodic Table of Dmitri Mendeleev (1834 -1907) By the late 1800s, chemists began to recognize that the elements can be arranged into groups or families based on similar chemical or physical properties. Dmitri Mendeleev, a Russian scientist, was credited for arranging the 60 elements known at that time into what is known today as the Periodic Table. Currently, about 116 elements are known in the Periodic Table. The Alkali Metals Several groups are known by their special names as a result of their characteristic chemical reactivities. The group 1A elements are also known as the ALKALI METALS. The word "alkali" is derived from an Arabic word meaning "ashes". Many sodium and postassium compounds were isolated from wood ashes (Na2CO3 and K2CO3 are still occasionally referred to as "soda ash" and "potash“, respectively). The alkali metals are: - soft, shiny metals - good conductors of heat and electricity - relatively low melting points - react vigorously with water - form white products when combined with Oxygen (Exception: Hydrogen, although listed as a group 1A element, is not a metal.) Examples of Alkali Metals 4 Li(s) + O2 (g) 2 Li2O(s) (lithium oxide, left) 2 Na(s) + O2 (g) Na2O2(s) (sodium peroxide, center) K(s) + O2 (g) KO2(s) (potassium superoxide, right) a) Lithium reacts with water to give off hydrogen gas (H2): Li(s) + 2 H2O(l) 2 Li(OH)(aq) + H2(g) b) Sodium reacts with water to give off hydrogen gas which can ignite spontaneously. c) Potassium reacts with water and ignites the evolved hydrogen gas very readily. The Group 2A elements (Be, Mg, Ca, Sr, Ba, Ra) are called the ALKALINE EARTH METALS. They are silvery-colored, The Alkaline soft and shiny metals like group 1A metals but not as reactive Earth Metals as the alkali metals. The reactivity of these elements increases going down the group. - they are found in compounds which form important rocks in the Earth's crust. This gives rise to the group's alternative name, the "alkaline earth" metals. The emerald (Be3(Al,Cr)2Si6O18) is the green variety of beryl, the mineral that is the principal source of beryllium. Radium is rare, white, luminescent, highly radio- active metallic element found in very small Emerald amounts in uranium ores. It is used in cancer radiotherapy, as a neutron source for some research purposes, and as a constituent of luminescent paints. Beryl A radium watch dial with radio- luminescent paints. Left image shows the areas of the watch that are the most radioactive (dark shadings). Solid Lumps of Beryllium metal. Beryllium has low aqueous solubility, which means it is rarely available to biological systems it has no known role in living organisms, and when encountered by them, is generally highly toxic. The Group 7A elements are called the HALOGENS, which include F, Cl, Br, I, and At. The halogens are strongly reactive, especially The Halogens fluorine and chlorine, and they form compounds with many other elements. In particular, the halogens form salt compounds with metals (e.g. NaCl, MgBr2, etc). - in their elemental forms, the halogens exist as diatomic molecules (e.g. F2, Cl2) but these only have a fleeting existence in nature and are much more common in the laboratory and in industry. At room temperature and pressure, fluorine and chlorine are gases, bromine is a liquid and iodine and astatine are solids; Group 7A is therefore the only periodic table group exhibiting all three states of matter. Bromine (Br2) Halogens are highly reactive, and as such can be harmful or lethal to biological organisms in sufficient quantities. Both chlorine and bromine are used as disinfectants for drinking water and swimming pools. Chlorine Chlorine is the active ingredient of most (Cl2) fabric bleaches and is used in the produc- tion of most paper products. Iodine (I2) The Noble Gases One of the eight main groups of elements, Group 8A is known as the Noble Gas group. Elements in this family have such a low reactivity that they were formerly known as the inert gases. Thus they are chemically very stable due to having the maximum number of valence electrons their outer shell can hold. The Noble Gases include helium, neon, argon, krypton, xenon, and radon. - under normal conditions, the noble gases occur as odorless, colorless, monatomic gases. Each of them has its melting and boiling point close together, so that only a small temperature range exists for each noble gas in which it is a liquid (e.g. Argon has mp: –189.6°C, bp: –185.8°C). The Noble gases have numerous important applications in lighting, welding and space technology. Argon is often used as a safe and inert Helium blimp atmosphere for the inside of filament light bulbs. Some of the noble gases glow distinctive colors Krypton laser when used inside lighting tubes (neon lights). Helium, due to its nonreactivity (compared with flammable hydrogen) and lightness, is often used in blimps and balloons. Krypton is also used in lasers, which are used by doctor for eye surgery. The Transition Metals The transition metals are the subgroups of elements intervening between groups IIA and IIIA in the periodic table. They are classified separately because of the filling of their d subshell orbitals. All the transition elements are metallic, but unlike the representative metals, they are likely to be hard, brittle, and have high melting points. Mercury (Hg) is an exception because it is a liquid at room temperature. - most have high electrical conductivity and can have a variety of oxidation states when forming compounds. Many transition metals are colored and as a result cause some of their ionic compounds to be colored. Naturally occurring copper: Copper occurs as the metal and as minerals such as blue azurite [2 CuO3·Cu(OH)2] and malachite [CuCO3·Cu(OH)2]. Small amounts of transition metal compounds are used to color glass: Cobalt oxide (Co2O3, blue); Copper or Chromiun oxides (green); Nickel oxides (purple); Uranium oxides (iridescent green). The inner transitional metals contain 2 main The Inner Transition Metals groups, the lanthanide series, and the actinide series. - like the Groups 1A and 2A elements, the inner transition metals are highly reactive. They catch fire in air easily, and react with water to liverate H2 lanthanides gas. Their reactivity makes them of very limited actinides use structurally. The lanthanide series include the 14 elements Gadolinium (Gd), a lanthanide metal, that proceed from atomic numbers 58 to 71. They is relatively stable in dry air. Its compounds are used in metallurgy, are shiny and silvery white metals, and tarnish optical instruments, and as easily when exposed to air. All lanthanides form contrasting agents in medical magnetic resonance imaging (MRI). MRI +3 ions as their principal chemical species. colonography The actinide series include the 14 elements that proceed from atomic numbers 90 to 103. These elements are unstable and their nuclei are prone to undergo radioactive decay. Unlike the lanthanides, the actinides show a variety of oxidation states. Uranium, for example, has compounds in each of the states, +3, +4, +5, and +6. Uranium (U) occurs naturally in soil and rock in trace amounts. When Uranium ore refined, it is silvery white and weakly radioactive. It is 70% denser than lead and is used in military high-density projectiles to destroy heavily armored targets. Glass containing uranium has a yellow green hue. Under UV light, uranium glass glows fluorescent green. Due to its radioactive nature, Uranium glass is used only in decorative items. Metals, Nonmetals, and Metalloids The heavy zig-zag line separates the elements into the metals and the nonmetals: - all elements on the left of the zig-zag line are metals (except Hydrogen). - the nonmetals are on the right of the zig-zag line. - the elements located along the zig-zag line are metalloids (B, Si, Ge, As, Sb, Te). Inner Transition Metals Some Characteristics of a Metal, a Nonmetal, and a Metalloid Silver (Ag) Antimony (Sb) Sulfur (S) Metal Metalloid Non-metal Shiny Blue-grey, shiny Dull, yellow Very ductile Brittle Brittle Malleable Shatters when Shatters when hammered hammered Good conductor of Poor conductor of Poor conductor, heat and electricity heat and electricity good insulator Density = 10.5 g/ml Density = 6.7 g/ml Density = 2.1 g/ml Melting point 962C Melting point 630C Melting point 113C Essential Elements The essential elements are vital to life that a deficiency in any one will result in either death, severe developmental abnormalities, or chronic ailments. - of the 116 elements known, 11 are predominant in biological systems. In humans, these 11 elements constitute 99.9% of the total number of atoms present, but 4 of these elements --- C, H, N, and O --- account for 99% of the total. The other 7 elements (Na, K, Ca, Mg, P, S, and Cl) comprise only 0.9% of the total atoms in the body; these generally occur as ions, such as Na+, K +, Mg2+, Ca2+, Cl–, and HPO42–. - the 11 essential elements are all “light” elements, having atomic numbers less than 21. Another 17 elements are also required by most by not all biological systems. These are generally “heavier” elements with atomic numbers greater than 18. They are about evenly divided between metals, nonmetals, and metalloids. Some Important Trace Elements In The Body Some metals and nonmetals known as trace elements are essential to the proper functioning of the body. Although they are required in very small amounts, their absence can disrupt major biological processes and cause illness. The trace elements listed below are present in the body combined with other elements. DV = Adult Daily Value Anemia, meaning "without blood", is a condition in Iron Deficiency Anemia which there is a deficiency of red blood cells (RBCs) and/or hemoglobin, the RBC protein that binds oxygen (O2) for transport to tissues. Without adequate RBCs or hemoglobin, blood has reduced ability to deliver O2 to local tissues, causing tissue hypoxia (low oxygen level in tissue). An anemic person often looks pale and fatigues easily. - a deficiency in body iron can lead to iron deficiency anemia (IDA), a condition in which blood lacks adequate healthy red blood cells. The body needs the element iron to make hemoglobin, an iron-rich substance that gives blood its red color and enables RBCs to carry O2 from the lungs to all parts of the body. In IDA, the RBCs are smaller than normal (microcytic) and have an increased zone of central pallor (hypochromic). Normal RBCs with biconcave shape and central pallor. Each RBC contains several hundred hemoglobin (Hb) molecules, each carrying four Fe atoms. RBCs from IDA patient are mirocytic and hypochromic. Symptoms of Iron Deficiency Anemia may include: - extreme fatigue, muscle weakness - pale skin, often cold hands and feet - shortness of breath, lightheadedness - headache - hair loss glossitis Other IDA symptoms may include: cheilosis - inflammation of the tongue (glossitis) koilonychia - brittle nails that are spoon-shaped (koilonychia) - scaling and fissure formation on the lips (cheilosis) - unusual cravings for non-nutritive substances such as ice, dirt, or pure starch (pica) - poor appetite, especially in infants and children with iron deficiency anemia IDA is common in 3 groups of population: 1) Pregnant women and growing children (esp. in developing countries). Increased body demands for iron in pregnant women and growing children are not being met with daily diet. 2) Menstruating women, especially those with menorrhagia (heavy monthly periods). Loss of iron due to heavy menstruation is not being compensated by dietary intake. 3) Older population. Chronic occult (not readily visible) blood loss can occur as a result of gastric ulcers, stomach and colon polyps, stomach and colon cancers, hemorrhoids, and others. It only takes about 1 to 2 teaspoons of blood loss daily to exceed iron absorption. Zinc Deficiency Leads to Acrodermatitis Enteropathica Zinc is an essential component of the diet. It is found in milk, meat, shellfish and wheat germ. Foods of plant origin are mostly low in zinc. Zinc is needed to assist metalloenzymes (proteins containing a metal atom) that are involved in many cellular processes throughout the body and in the normal functioning of the brain. A deficiency in dietary zinc can lead to the following characteristic manifestations: - skin inflammation (dermatitis) that evolve into crusted, blistered, pus-filled and eroded lesions. This occurs particularly around body openings such as the mouth, anus, and eyes, and the skin on elbows, knees, hands, and feet. - diffuse hair loss on the scalp, eyebrows and, eyelashes. - impaired wound healing; secondary bacterial infection. - in severe cases, irritability and emotional disturbances are evident due to wasting (atrophy) of the brain cortex. Treatment: Supplemental zinc usually eliminates the symptoms. Clinical Case: a 6-week-old male infant developed a widespread red crusted and eroded dermatitis that was most prominent on the distal extremities, creases, and around body orifices. He was listless and had frequent watery stools. His skin lesions and systemic symptoms cleared within several days of initiating high dose oral zinc supplementation. Zinc levels obtained at the time of admission to the hospital were extremely low. Dalton’s Atomic Theory All the elements in the Periodic Table are made up of atoms, which are the smallest particles of matter that retain the characteristics of the elements. - in the 19th century, John Dalton developed the Atomic Theory, propounding the idea that atoms are responsible for the combinations of elements found in compounds. The Dalton's Atomic Theory has the following postulates: 1. All matter is made up of tiny particles called atoms. Example: Salt is made up of Sodium atoms and Chlorine atoms. 2. All atoms of a given element are similar to one another and different from atoms of a different element. 3. Atoms of 2 or more different elements combine to form compounds. . A particular compound is always made up of the same kinds of atoms . that are in the same proportions. Ex: Water is made up of 1 Oxygen atom and 2 Hydrogen atoms. 4. A chemical reaction involves rearrangement, separation, or combination of atoms. . Atoms are never created nor destroyed in a chemical reaction. The total number of atoms for each element 4 Fe + 3 O2 2 Fe2O3 is preserved in a chemical reaction. Particles in the Atom In the early 20th century, growing experimental evidence indicated that atoms are not solid spheres as scientists have imagined but composed of smaller particles called subatomic particles. ● 3 subatomic particles are of interest to us: i. Proton has a (+) charge and a mass = 1.7 x 10–24 g. ii. Neutron has no electrical charge and is therefore neutral; it has the same mass as the proton. iii. Electron has a (–) charge; its mass is about 1/2000 the mass of either the proton or neutron and is usually ignored in atomic mass calculations. Subatomic particles of like charges repel one another, whereas particles of opposite charges attract one When dry hair is brushed another. with a balloon on a dry day, it is attracted to the - electrical charges have been balloon. noted early in the development of the theory of the atom. Examples of electric charge phenomena that we see in everyday life are shown Rub a glass rod with a silk cloth, at right. it will attract a pile of confetti. Because the value 1.7 x 10–24 g is at such a tiny scale, scientists defined another convention for measuring the mass of matter at the atomic level – the atomic mass unit (amu). Definition: One atomic mass unit (1 amu) is defined as one-twelfth of the mass of an atom of carbon with 6 protons and 6 neutrons. - such a carbon atom is assigned an exact mass of 12.000 amu and is called Carbon-12. Masses of other atoms are assigned relative to the mass of carbon-12. Ex: an Oxygen atom with 8 protons and 8 neutrons has a relative mass of 16.000. The atomic mass (A) of an atom can be estimated by summing its number of protons and neutrons: A = number of protons + number of neutrons Ex: a sodium atom (Na) with 11 protons and 12 neutrons has a mass of A = 23. The amu is also called dalton in honor of the 19th century chemist John Dalton. On the atomic scale: mass of a proton = mass of a neutron = 1 amu = Summary of Subatomic Particles Subatomic Symbol Electrical Mass Location particle charge (amu) in atom Proton p or p+ +1 1 Nucleus Neutron n or n0 0 1 Nucleus Electron e– –1 1/2000 Outside nucleus - the mass of a proton is ~1.7 x 10–24 g - 1 amu = 1/12 the mass of carbon-12 (which has 6 protons and 6 neutrons). - on the amu scale: mass of proton = mass of neutron. - 1 amu is also known as 1 dalton. - mass of an electron is usually ignored. Rutherford’s Gold Foil Bombardment Experiment While various experiments have indicated the existence of subatomic particles such as the electron and the positively charged particles, scientists were not certain how these particles are arranged in an atom. Further, if electrons are very much lighter than atoms, then the positively charged particles must carry the mass of the atom. JJ Thomson suggested that atoms are spheres of positive charge in which light, negatively charged electrons are embedded, much as raisins might be embedded in the surface of a pudding. This came to be known as the Raisin Pudding Model of the Atom. ● In the early 1900s, Ernest Rutherford performed a gold foil bombardment experiment using a beam of positively charged particles that have the same mass as the Helium nucleus (e.g. particles). Where he expected the particles to pass undeflected through the gold foil, to his surprise he found a few were deflected at various angles. Some even came backwards. Rutherford’s Gold Foil Experiment. A beam of (+)charged particles was directed at a thin gold foil. A fluorescent screen was used to detect Ernest Rutherford particles passing through. Most particles passed through the foil, but some were deflected from their path. A few were even deflected backward. Rutherford reasoned that only a very concentrated positive charge in a tiny space within the gold atom could possibly repel the fast-moving, particles enough to reverse the particles’ direction. He hypothesized that the mass of this positively charged region, which he called the nucleus, must be larger than the mass of the particle. He went on to argue that the reason most of the particles were undeflected was that most parts of the atoms in the gold foil were empty space. Top: Expected results of Rutherford's gold foil experiment: Alpha particles passing through the raisin pudding model of the atom undisturbed. Bottom: Observed results: Some of the particles were deflected, and some by very large angles. Rutherford concluded that the positive charge of the atom must be concentrated into a very small location: the atomic nucleus. A New Model Of The Atom From the results of the gold foil experiment, Rutherford proposed a new model of the atom in which all of the positive charges and most of the mass of the atom are concentrated in a very small core, the nucleus. The electrons occupy the rest of the space in the atom. Most of the atom is actually “empty space”. Structure Of The Atom As a result of Rutherford’s gold foil experiment (and other scientists’), the general structure of the atom has come to accepted as having a small, dense core region known as the nucleus, where the subatomic particles protons and neutrons are located. The nucleus is situated at the center of the atom and has a positive charge due to the presence of protons. - most of the rest of the atom is empty space, occupied by fast moving electrons. Atomic Number and Mass Number An element is defined by the number of protons in its nucleus. It is the number of protons that give an element its characteristics and distinguish it from another element. The number of protons in the nucleus of an atom is also referred to as the atomic number, which is the whole number written above the atomic symbol in the Periodic Table. ● The Periodic Table arranges the elements in order of increasing atomic number. Example: Lithium has an atomic number = 3 lithium has 3 protons Carbon has an atomic number = 6 carbon has 6 protons ● An atom is electrically neutral, that is, all the (+) charges of the protons in the nucleus are balanced by an equal number of (–) charges in the electrons. Thus, there are as many electrons in a neutral atom as there are protons. - in other words, in a neutral atom, the atomic number, which is the number of protons in an atom, is also the number of electrons. Atomic Mass: The mass of an atom is the sum of its numbers of protons and neutrons: Atomic Mass = # Protons + # Neutrons Examples: - Oxygen (O) atom has 8 protons and 8 neutrons atomic mass = 16 - Iron (Fe) atom has 26 protons and 30 neutrons atomic mass = 56 - Gold (Au) atom has 79 protons and 118 neutrons atomic mass = 197 Atoms and their Subatomic Particles Atomic Mass Number Number Number Element Symbol number number of protons of neutrons of electrons Hydrogen H 1 1 1 0 1 Nitrogen N 7 14 7 7 7 Chlorine Cl 17 37 17 20 17 Iron Fe 26 56 26 30 26 Gold Au 79 197 79 118 79 - note that the atomic number and the number of electrons are identical. - the mass number is the sum of the numbers of protons and neutrons. - atoms with increasingly higher atomic mass tend to have increasingly higher number of neutrons relative to the number of protons. Number of protons ? neutrons ? electrons ? An atom of phosphorus (P) has an atomic mass = 31 - what are the number of protons ? - what are the number of electrons ? - what are the number of neutrons ? Would you use atomic number, atomic mass, or both to obtain the following ? - the number of protons in an atom ? - the number of neutrons in an atom ? - the number of electrons in a neutral atom ? - the number of particles in the nucleus ? Writing Atomic Symbol An atom is written with a convention that gives information on its atomic number and mass number: 24 top number indicates mass number (# protons + # neutrons) Mg 12 bottom number indicates atomic number (# protons) - note that this writing convention is not the same way as the Periodic Table presents the elements. In the Periodic Table, the top number represents the number of protons instead of the mass, but this is the exception that is found in the Periodic Table. In the Periodic Table, the atomic number is commonly listed as a superscript or at the top of the symbol of the element (and the mass number at the bottom). Isotopes It was found that not all atoms of the same element have exactly the same mass number. That is, if you can examine a sample of most any element at the atomic level, you will find that while the majority of the atoms in the sample has the same mass number, a small percentages of these atoms have slightly different mass numbers. Ex: Magnesium (atomic number 12) can have 3 slightly different masses that differ in their number of neutrons: 24 25 26 Mg (79%) Mg (10%) Mg (11%) 12 12 12 - all atoms of the element Magnesium has the same number of protons (12 protons). The percentages given in parentheses indicate the average natural abundances of these atoms in a pure sample of magnesium. - however, some has: 12 neutrons ( to give atomic mass 24) 13 neutrons ( to give atomic mass 25) 14 neutrons ( to give atomic mass 26) - we called atoms of the same element that have different masses Purified magnesium is a silvery, white metal. By due to variable numbers of neutrons isotopes. While isotopes of mass, it makes up 2% of the Earth’s crust and and the the same element have different masses, their chemical behavior 4th most abudant element in is very similar. our body. Examples of Isotopes Atoms that have the same number of protons but different numbers of neutrons are called isotopes. The element Hydrogen, for example, has three commonly known isotopes ─ protium (1H), deuterium (2H), and tritium (3H). Helium and Lithium each has two stable isotopes. - some isotopes are classified as stable isotopes and others are classified as unstable, or radioactive, isotopes. Stable isotopes maintain constant concentrations on Earth over time. - by contrast, unstable isotopes are atoms that disintegrate (or decay) at predictable and measurable rates to form other isotopes by emitting a nuclear electron or a helium nucleus and radiation. These isotopes continue to decay until they reach stability. As a rule, the heavier an isotope is than the most common isotope of a particular element, the more unstable it is and the faster it will decay. Average Atomic Mass (Weight) The percentages of the different isotopes of an element are experimentally determined. Because a sample of an element often has a mixture of 2 or more isotopes of different abundances, the mass of an element as shown on the Periodic Table is actually the average of all the masses of the isotopes that are found for the given element. This is one reason why the atomic masses listed on the Periodic Table are not a whole number. Ex: Calculate for the average mass of Magnesium (Mg) based on their isotope abundances (magnesium-24: 79%; magnesium-25: 10%; magnesium-26: 11%). Solution: Weighted average mass: 24 (0.79) + 25 (0.10) + 26 (0.11) = 24.32 - the weighted average mass is calculated by multiplying the mass of each isotope by its fractional abundance and summing all the isotopic masses. Because the average mass 24.32 is closer to isotope 24 than isotope 25 or 26, Mg-24 isotope is the most prevalent in a magnesium sample. Atomic % Natural Isotope Mass Abundance Isotopes of Neon. Based on natural abundances, 20Ne 19.99 90.9% the average atomic mass of neon is closest to 21Ne 20.99 0.3% which whole number ? 22Ne 21.99 8.8% Practice Problems 1. Copper contains two isotopes, copper-63 and copper-65. Are there more atoms of copper-63 or copper-65 in a sample of copper ? 2. What are the number of protons, neutrons, and electrons in the following isotopes ? Aluminum metal rods 27 106 Al Cd 13 48 3. Write the atomic symbols for isotopes with the following: i. 26 electrons and 30 neutrons Pure cadmium metal is soft, malleable, ductile, and has a ii. a mass number 24 and 13 neutrons white bluish color. It is similar to zinc. We learn that an atom consists of a central nucleus made of protons and neutrons, and about which revolve electrons like planets of n=2 the solar system revolve about the sun. n=1 - however, electrons don't just move about randomly within the space around the nucleus, Niels Bohr’s Model of the but rather there are distinct levels of energy in which electrons with certain energies occupy Atom Nucleus and move about. - Niels Bohr was one of the first physicists early in the 20th century who propounded the idea of quanta of energy to explain the atomic line spectrum of the Hydrogen atom. According to Niels Bohr, there are energy levels called principal quantum energy levels (also called main energy levels, or energy shells), conventionally given the symbol n, that are closer to the nucleus in which electrons with lower energy occupy, and there are quantum energy levels that are farther away from the nucleus for electrons with higher energies to occupy. The above diagram shows an atom with the first two principal energy levels n = 1 and n = 2. - for an electron to move from a lower energy level to the next higher level requires that it gains or absorbs energy. Conversely, an electron drops from a higher energy level to a lower one by losing energy so that it now moves closer to the nucleus. We can think of the quantum energy levels of an atom as similar to the rungs on a ladder. The lowest energy level would be the 1st rung, the second energy level the 2nd rung, and so on. As one climbs up or down the ladder, one must step from one rung to the next and can not stop at a level between the rungs. - similarly, electrons in atoms exist in the available energy levels. To move up to the next higher energy level (n+1), an electron must absorb a definite quantum of energy (E = h). Conversely, to fall to E = h the next lower energy level (n–1), an electron must lose the same amount of quantum of energy. - unlike the ladder, the lower energy levels in an atom are far apart compared to the higher energy levels that are closer together. The maximum number of electrons allowed in each energy level is given by the formula: 2n2 Using this formula, one can calculate the total number of electrons in any given energy level as shown in the Table at right. The lowest energy level (n = 1) can hold up to 2 electrons; energy level n = 2 can accommodate up to 8 electrons; etc. In the atoms of the elements known today, electrons occupy 7 different energy levels. When atoms of an element are energized, such as in the red glow of a metal strip produced by a hot flame or the white radiance of a bulb filament produced by electrical energy, the electrons in those atoms absorb a certain amount of energy and jump to higher energy levels. When some of the energized electrons drop back to lower, more stable energy levels, energy is emitted in the form of radiation. The “radiation rays” emitted by energized atoms can have different wavelengths (often measured in nm). If these radiation wavelengths fall within the visible range of the electromagnetic spectrum, our eyes can detect them as visible light emission. Depending on the wavelengths of radiation, the light emission takes on the characteristic color produced by the energized atoms of the element. Radiation within the visible region of the electro- magnetic spectrum produced by hydrogen gas. Line Spectrum of Hydrogen Line Spectrum of Hydrogen When H2 gas is excited by an electrical current, it emits radiation that we can see as a glow of light. This light is actually a mixture of light energy emitted at different wavelengths. When this light is passed through a glass prism, it is resolved into four discrete bands of light, each having a different wavelength and associated color. Together, the bands of light are called the visible line spectrum of hydrogen. Sunlight is diffracted in a rainbow of colors in a glass prism. A glass prism resolves light emitted by energized hydrogen gas into four discrete bands, each with a different wavelength and color. Orbitals and Electron Arrangements in Atoms The principal quantum energy levels (or energy shells) are equivalent to the Periods of the Periodic Table. Ex: Shell n = 1 corresponds to Period 1, which contains the elements Hydrogen and Helium and holds a maximum of 2 electrons (in He). - similarly, shell n = 2 corresponds to Period 2, which contains the eight elements Li, Be, B, C, N, O, F, and Ne and holds a maximum of 8 electrons (in Ne). Electron Orbitals The view of electrons orbiting in energy shells around the nucleus of an atom as planets of the solar system orbiting the sun is a simplistic model of the atom. An implied assumption of this model is that electrons belonging to the same energy shell all have equal energy. However, early in the 20th century, with the development of quantum mechanics – a series of complex mathematical equations that describe the position of an electron around a nucleus – scientists began to realize that electrons belonging to a given energy shell can have different energies and, depending partly on their energies, tend to occupy certain spatial regions with peculiar shapes called orbitals. Electron Orbitals. An orbital is defined as a region in space around the nucleus in which electrons is most likely to be found. Each orbital can hold a maximum of 2 electrons. There are several types of orbitals, of which four are of interest to us – the s, p, d, and f orbitals. - an s orbital is spherical with the nucleus at the center. - a p orbital has 2 lobes, each lobe being the region of space having the highest probability of finding an electron. The 2 lobes ended at the node where the nucleus is located. There is a set of three p orbitals arranged in the three directions x, y, and z around the nucleus, forming px, py, and pz orbitals. These three p orbitals can accommodate up to 6 electrons. - there are a set of five d orbitals and seven f orbitals; however, their geometry is much too complex for our purpose. The five d orbitals can house up to 10 electrons, and the seven f orbitals up to 14 electrons. Divisions of the s-, p-, d-, and f-Orbitals in the Periodic Table The lowest energy shell (n = 1) consists only of a single s orbital. Shell n = 2 can have both s and p orbitals. Shell n = 3 can have s, p, and d orbitals. Higher energy shells (n = 4 and above) can have s, p, d, and f orbitals. The distribution of the orbitals based on electron arrangements of elements in the Periodic Table is given in the chart below. Energy shell n=1 n=2 n=3 n=4 n=5 n=6 n=7 Orbitals in Energy Shells n = 1 to n = 4 Orbital distribution for the first 4 energy shells is given the Table below. Recall that the maximum number of electrons that can be found in a given shell is given by the formula 2n2. Electron Level Arrangements for Elements in the First 4 Shells The electron level arrangement of an atom gives the number of electrons in each energy level. When a shell is completely filled to its maximum allowed number of electrons, additional electrons are placed in the next higher shell. Number of electrons Atomic in Energy Shell Element Symbol number 1 2 3 4 hydrogen H 1 1 helium He 2 2 (1st shell completed) lithium Li 3 2 1 beryllium Be 4 2 2 boron B 5 2 3 neon Ne 10 2 8 (2nd shell completed) sodium Na 11 2 8 1 magnesium Mg 12 2 8 2 aluminum Al 13 2 8 3 argon Ar 18 2 8 8 potassium K 19 2 8 8 1 calcium Ca 20 2 8 8 2 scandium Sc 21 2 8 9 2 zinc Zn 30 2 8 18 2 (3rd shell completed) gallium Ga 31 2 8 18 3 krypton Kr 36 2 8 18 8 Periodic Table The chemical properties of representative elements are mostly due to their valence electrons, which are the electrons Valence Electrons in the outermost energy level. The valence electrons are the electrons that participate in chemical reactions. - in Chem109, we define the valence electrons as the s and p electrons. The Group Numbers indicate the valence (outer) electrons for the elements in each vertical columns. Ex: Group 1A elements (e.g. Li, Na, K, etc) all have 1 valence electron in the outer energy shell. Group 2A elements (e.g. Be, Mg, Ca, etc) all have 2 valence electrons in the outer shell. Group 7A elements (e.g. F, Cl, Br, etc) all have 7 valence electrons in the outer shell. Electron-Dot Symbols and the Octet Configuration The valence electrons are represented as dots placed along the 4 edges of the symbol of the element. When there are more than 4 electrons, the electrons are paired up. For any given shell, the valence electrons fill up the s and p orbitals until reaching the Noble Gas configuration with 8 valence electrons. The atom is said to achieve the Octet configuration and chemically becomes very stable. For this reason, the Noble Gases are unreactive and usually do not form compounds with other elements. Trends in Atomic Sizes for the Representative Elements Atomic size can be taken approximately as the distance from the nucleus to the valence electrons on the outermost shell. radius - atomic size tends to increase going down a group. Each successive increase in the energy level puts the outer electrons further away from the nucleus. - atomic size typically decreases from left to right across a period. Going across a period, as the number of valence electrons increases, the number of protons in the nucleus also increases, thereby pulling the valence electrons closer to the nucleus. This makes the atom shrinking smaller in size. First Ionization Energy of the Representative Elements In a neutral atom, the (–)charged e– move about the (+)charged central nucleus. The neutral atom can acquire a net (+) charge when an e– is removed from the atom. An atom that has a positive charge is called a cation. To pull an e– away from the (+)charged nucleus requires energy. The ionization energy is defined as the energy needed to remove an electron from an atom in the gaseous state. Ex: Na(g) + energy Na+(g) + e– The ionization energy generally The ionization energy generally increases going decreases going down a group across a period from left to right because as the because less energy is needed to number of protons increases, the nuclear attraction remove an e– as nuclear attraction for for e– becomes increasingly stronger. e– decreases farther from the nucleus.