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Group 2 Project
Transcript of Group 2 Project
Neutrons: 20 Bright Line spectrum Physical Properties Density (near r.t.):1.55 g•cm−3 Liquid density at m.p. : 1.378 g•cm−3 Melting point:1115 K, 842 °C, 1548 °F Boiling point: 1757 K, 1484 °C, 2703 °F Heat of fusion: 8.54 kJ•mol−1 Heat of vaporization:154.7 kJ•mol−1 Solid at room temperature
Calcium is a soft gray alkaline earth metal Chemical Properties The metal is rather reactive. It readily forms a white coating of calcium nitride (Ca3N2) in air. It reacts with water and the metal burns with a yellow-red flame, forming largely the nitride. Calcium reacting with water. Calcium-48 P:20
N: 28 Abundace
.187% HL: 4.3×1019 yrs Radioactive Calcium-46 P: 20
.004% Stable Isotopes Calcium-44 Stable P:20
2.086% Calcium-43 P:20
.135% Stable Calcium-42 P: 20
N: 22 Abundance
.647% Stable Calcium-40 P:20
N:20 Stable Abundance
96.964% Barium Protons:56
Electrons:56 Appearance Discovery Calcium was discovered by Sir Humphrey Davy at 1808 in England.
Origin of name: from the Latin word "calx" meaning "lime".
Calcium metal was not isolated until 1808. After learning that Berzelius and Pontin prepared calcium amalgam by electrolysing lime in mercury,
Sir Humphry Davy was able to isolate the impure metal.
He did this by the electrolysis of a mixture of lime and mercuric oxide (HgO). Calcium metal was not available in large scale until the beginning of the 20th century. Sir Humphrey Davy Berzelius What is electrolysis? is a process in which electrical energy is used to produce a chemical change that would not otherwise occur. First, an electric current is produced from a chemical using a oxidation-reduction reaction. Then this electric current is used to produce a chemical change. An example of electrolysis is hydrolysis, or the splitting of water into hydrogen and oxygen. Where does Calcium come from? Calcium does not occur as the metal itself in nature and instead is found in various minerals including as limestone, gypsum and fluorite. Calcium metal is readily available commercially and there is no need to make it in the laboratory. Commercially it can be made by the electrolysis of molten calcium chloride, CaCl2.
cathode: Ca2+(l) + 2e- → Ca
anode: Cl-(l) → 1/2Cl2 (g) + e- Limestone Gypsum Fluorite Uses of Calcium Humans: Medicine: Calcium Suppplements-This medication is used to prevent or treat low blood calcium levels in people who do not get enough calcium from their diets. It may be used to treat conditions caused by low calcium levels such as bone loss (osteoporosis), weak bones (osteomalacia/rickets), decreased activity of the parathyroid gland (hypoparathyroidism), and a certain muscle disease (latent tetany) Industry: Chemical industry, calcium oxide is used as a dehydrating agent and an absorbent. The former is important for purifying other compounds and substances such as glucose, citric acid and certain dyes before the compound undergoes further refinement. As an absorbent, quicklime has been shown to effectively absorb carbon dioxide fumes, which helps when the chemical’s carbon dioxide is too high to effectively work with. Abundance In Humans: •Ninety-nine percent of the calcium in the human body is stored in the bones and teeth. The remaining one percent is found in bones and other tissues. In Earth: In Universe: Hazards To Humans: Constipation is a common adverse reaction to getting too much calcium. Stomach upset, nausea, vomiting, increased urination, confusion, loss of appetite and an irregular heart rhythm may also occur
To the Environment: Calcium phosphide is very toxic to aquatic organisms Agriculture: Growing a garden and raising animals, including chickens, are important practical activities in this course. Calcium relates to soil fertility and to the ability of chickens to produce eggshells. Also, since humans can't move a muscle without the presence of calcium in the muscle, it relates to the students' ability to do agricultural work. Interesting Facts Have you ever wanted to be ‘in the limelight?’ Lime is calcium oxide, which produces a brilliant, intense light when burnt in an oxyhydrogen flame. It was used to light the stage in theaters during the 1800s until electricity took over Calcium metal burns in air, forming a mixture of calcium oxide and nitride. Cells in animals and plants must communicate with other cells. This is called signaling. Calcium ions are the most important messengers between cells in living things and are absolutely vital for the existence of multicellular life forms. Strontium Bohr Diagram P: 38
E: 38 Bright Line Spectrum Physical Properties Melting point: 1050 [or 777 °C (1431 °F)] K
Boiling point: 1655 [or 1382 °C (2520 °F)] K
Strontium is a soft, silvery metal. When cut it quickly turns a yellowish color due to the formation of strontium oxide (strontia, SrO). Finely powdered strontium metal is sufficiently reactive to ignite spontaneously in air.
Solid at room temperature. +2 Chemical Properties More reactive in water than calcium,
It reacts on contact to produce strontium hydroxide and hydrogen gas.
It burns in air to produce both strontium oxide and strontium nitride, but since it does not react with nitrogen below 380 °C, at room temperature it will only form the oxide spontaneously Strontium Powder Strontium metal and water Isotopes Strontium-88 Stable Abundance
E: 38 Strontium-87 Stable Abundance
E:38 Strontium-86 Stable Abundance
E:38 Strontium-84 Stable Abundance
E: 38 Strontium has two radioactive isotopes but they have a 0% abundance. They are Strontium-89 (HL: 50.53 d) and Strontium 90 (HL: 28.79 yrs). Appearance Freshly cut strontium has a silvery appearance, but rapidly turns a yellowish color with the formation of the oxide. Discovery Strontium was recognized as a new element in 1790 by Adair Crawford
analyzed a mineral sample from a lead mine near Strontian, Scotland.
Strontium metal was first isolated in 1808 by Sir Humphry Davy working in London using electrolysis. How did he do it? Davy had built a very large 600-plate battery, which he used to pass electricity through salts, breaking them down in order to isolate new elements.
He mixed magnesium oxide to a paste with (probably) strontium sulfate.
He made a depression in the paste and placed mercury metal there to act as an electrode. Platinum was used as a counter electrode
When electricity was passed through the paste, a strontium-mercury amalgam formed at the mercury electrode. Where does Strontium come from? Strontium is never found free in nature. The principal strontium ores are celestine (strontium sulfate, SrSO4) and strontianite (strontium carbonate, SrCO3).
The main commercial process for strontium metal production is reduction of strontium oxide with aluminum. What can Sr be used for? Humans/Medicine:
Strontium chloride is used in toothpaste for sensitive teeth.
strontium was found to have anabolic activity in bone, and thus may have significant beneficial effects on bone balance in normal and osteopenic animals.
Strontium has been thought to have potential in the treatment of osteoporosis.
Strontium 89 (Metastron). Strontium is a chemical that tends to collect in bones. This is because strontium is very similar in structure to calcium, which bones are mostly made of. And strontium 89 is a radioactive form of strontium. Industry:
Strontium is used for producing glass (cathode ray tubes) for color televisions.
It is also used in producing ferrite ceramic magnets and in refining zinc.
Fire Works Cathode ray tube Abundances Abundance earth’s crust: 370 parts per million by weight, 87 parts per million by moles
Abundance solar system: 50 parts per billion by weight, 0.7 parts per billion by moles
Abundance human body: About 99% of the strontium in the human body is concentrated in the bones Hazards Humans:
Strontium-90 (90Sr) has a half-life of 28 years. It is a product of nuclear fallout and presents a major health problem
Strontium-89, 90 takes on the characteristics of calcium which the body directs to the bones and teeth and is considered one of the more hazardous constituents of nuclear wastes causing leukemia and bone caner (EPA). Environment:
Strontium is always present in air as dust, up to a certain level.
One of the isotopes of strontium is radioactive. This isotope is not likely to occur naturally in the environment.
All strontium will eventually end up in soils or bottoms of surface waters, where they mix with strontium that is already present. Protons: 4
Electrons: 4 Protons: 12
Electrons: 12 Interesting Facts:
The world’s most accurate atomic clock, accurate to one second in 200 million years, has been developed using strontium atoms. Moon dust mainly consists of strontium aluminate. This same compound is also found in phosphorescence materials that glow in the dark, such as luminous paints, party-wear and toys. Be-7 P: 4
N: 3 Half-life:
53.22 d Abundance:
Trace amounts Be-10 P: 4
N: 6 Half-life:
1.51x10^6 y Abundance:
Trace amounts Be-9 Bibliography:
Winter, Mark. "Calcium." WebElements Periodic Table of the Elements. The Univeristy of Sheffield, n.d. Web. 13 Jan. 2013.
"Calcium Element Facts." Chemicool. N.p., n.d. Web. 13 Jan. 2013.
Johnson, Steve. "Physical or Unusual Facts About Strontium." EHow. Demand Media, 09 May 2010. Web. 13 Jan. 2013.
Dayah, Michael. "Dynamic Periodic Table." Dynamic Periodic Table. N.p., n.d. Web. 13 Jan. 2013.
"Beryllium." Chemicool Periodic Table. Chemicool.com. 15 Oct. 2012. Web. 1/15/2013
"Magnesium." Chemicool Periodic Table. Chemicool.com. 07 Oct. 2012. Web. 1/15/2013
"Dynamic Periodic Table." Dynamic Periodic Table. N.p., n.d. Web. 15 Jan. 2013.
"Chemistry Explained." Beryllium, Chemical Element. N.p., 2013. Web. 15 Jan. 2013.
Winter, Mark J. "Beryllium." WebElements Periodic Table of the Elements. N.p., n.d. Web. 15 Jan. 2013.
Bentor, Yinon. "Chemical Elements.com - Radium (Ra)." Chemical Elements.com - Radium (Ra). N.p., n.d. Web. 15 Jan. 2013.
Winter, Mark. "Radium." WebElements Periodic Table of the Elements. N.p., n.d. Web. 15 Jan. 2013.
"Radium Element Facts." Chemicool. N.p., n.d. Web. 15 Jan. 2013.
Bentor, Yinon. "Chemical Elements.com - Barium (Ba)." Chemical Elements.com - Barium (Ba). N.p., n.d. Web. 15 Jan. 2013.
Winter, Mark. "Barium." WebElements Periodic Table of the Elements. N.p., n.d. Web. 15 Jan. 2013.
"Barium Element Facts." Chemicool. N.p., n.d. Web. 15 Jan. 2013. Why are these elements grouped together?
Similar Physical and Chemical Characteristics
All have 2 valence electrons
All for +2 ions
Found in Compounds
All metals with a shiny, silvery-white color ` P: 4
N: 5 Abundance:
About 100% Stable Isotopes Density (near r.t.): 1.85 g·cm−3
Liquid density at m.p.: 1.690 g·cm−3
Melting point: 1560 K, 1287 °C, 2349 °F
Boiling point: 2742 K, 2469 °C, 4476 °F
Heat of fusion: 12.2 kJ·mol−1
Heat of vaporization: 297 kJ·mol−1 Physical Properties Steel gray color
Solid at STP, brittle at room temperature
Has high thermal conductivity and is nonmagnetic
Crystal structure: hexagonal Chemical Properties - An oxide layer forms around beryllium that prevents further reactions with air unless heated above 1000°C.
-When ignited, burns brilliantly forming a mixture of beryllium oxide and beryllium nitride.
-Dissolves in non-oxidizing acids, such as HCl and diluted H SO , and in alkali solutions, but not in nitric acid or water because this forms the oxide layer.
-This behavior is similar to that of aluminium metal. 2 4 Discovery -First discovered in 1798, in France, by René Haüy who saw similarities in the crystal structures and properties of beryl and emerald and wondered if there was a common element in both.
-Hauy asked Nicolas Louis Vauquelin, a French chemist who specialized in analysis, to see what he could find and Vauquelin discovered a new, sweet-tasting substance in both emerald and beryl that had been previously overlooked due to its similarity to aluminum. We now call this substance beryllia, BeO, but originally the sweet taste of the salts led to the new element being named ‘glyceynum,’ then ‘glucina’ or ‘glucine.’ Despite its sweet taste, we now know that beryllium and its compounds are highly toxic.
-In 1957, Beryllium became the element’s official name.
-Friedrich Wöhler (German) and Antoine Bussy (France) independently isolated beryllium in 1828 by the chemical reaction of metallic potassium with beryllium chloride. Sources:
-The mineral beryl, Be Al (SiO )
-The two main ores of beryllium, beryl and bertrandite, are found in Argentina, Brazil, India, Madagascar, Russia and the United States. Total world reserves of beryllium ore are greater than 400,000 tonnes
-Commercially it is produced by the
reduction of the fluoride with magnesium metal. 3 2 3 6 Uses:
-Is virtually transparent to x-rays and is used in radiation windows for x-ray tubes.
-Beryllium alloys are used in the aerospace and defense industry as light-weight materials for high performance aircraft, satellites and spacecraft, and guided missiles.
-Beryllium is used as an alloy with copper to make spark-proof tools.
-Beryllium is also used in nuclear reactors as a reflector and absorber of neutrons, a shield and a moderator. Hazards:
its salts are both
toxic and carcinogenic. Fun Facts:
-Beryllium played a large part in proving the existence of neutrons. In 1932, James Chadwick, an English physicist, bombarded a sample of beryllium with alpha-rays (helium nuclei). He observed that the bombarded sample emitted a subatomic particle, which had mass but no charge.
This neutral particle was the neutron
Though beryllium tastes sweet, it is toxic. Chronic exposure to beryllium (typically through inhalation) can lead to a life-threatening allergic disease called berylliosis. Abundance:
Because any beryllium synthesized in stars is short-lived, it is a relatively rare element in both the universe and in the crust of the Earth
-Earth's crust: 2.8 parts per million by weight
-Total world reserves of beryllium ore are greater than 400,000 tonnes. Mg-24 Stable Abundance:
78.99% P: 12
N: 12 Mg-25 P: 12
N: 13 Abundance:
10% Stable Mg-26 Stable Abundance:
11.01% P: 12
N: 14 Isotopes Physical Properties -Silvery-white
-Solid at STP
-Fairly strong, light weight metal
-Crystalline structure: hexagonal close-packed Chemical Properties -Highly flammable metal when powdered or shaved into thin strips
-Once ignited, it is difficult to extinguish, being able to burn in nitrogen, carbon dioxide and water.
-when it reacts with hydrochloric acid to produce the chloride of the metal and releases hydrogen gas
- magnesium reacts with water at room temperature, like calcium, though it reacts much more slowly.
-Magnesium's ability to react with water can be harnessed to produce energy and run a magnesium-based engine.
-tarnishes in air to form a thin oxide coating.
-Mg and its alloys have very good corrosion resistance.
-When it burns in air, produces a brilliant white light, and produces magnesium oxide and magnesium nitride. Density (near r.t.): 1.85 g·cm−3
Liquid density at m.p.: 1.690 g·cm−3
Melting point: 1560 K, 1287 °C, 2349 °F
Boiling point: 2742 K, 2469 °C, 4476 °F
Heat of fusion: 12.2 kJ·mol−1
Heat of vaporization: 297 kJ·mol−1 -Magnesium and calcium were once thought to be the same substance.
-1755: Scottish chemist, Joseph Black, showed by experiment that the two were different.
-1808: First isolated by Sir Humphrey Davy, in London, England. Davy had built a large battery and used it to pass electricity through salts using electrolysis of a mixture of magnesia and mercuric oxide. In doing so, he discovered or isolated for the first time several alkali and alkali earth metals, such as magnesium, calcium, and strontium.
-1830: France, Antoine Bussy published his work showing how pure magnesium metal could be obtained. Discovery Sources: USES OF MAGNESIUM
-Used in iron and aluminum and structural metal
-Electronic devices: mobile phones, laptop and tablet computers, cameras, and other electronic components.
The brilliant light it produces when ignited is made use of in photography, flares, pyrotechnics and incendiary bombs.
Due to very low density, magnesium alloys are used in aircraft, car engine casings, and missile construction.
Some magnesium compounds are used for medicinal purposes. Needed in many bodily functions such as:
-the proper growth, formation and function of our bones and muscles, prevents some heart disorders and high blood pressure, metabolism, helps our brains function normally, control insulin levels in your blood. Humans: Adult human bodies contain about 24 grams of magnesium, with 60% in the skeleton, 39% intracellular (20% in skeletal muscle), and 1% extracellular
Earth's crust: 2.3 % by weight
-Magnesium is the eighth most abundant element in the Earth’s crust by mass and the sixth most abundant metal
Solar system: 700 parts per million by weight Harmful effects:
is an explosive hazard.
bright white light plus
ultraviolet from burning magnesium
can cause permanent eye damage. Fun Facts: -First isolated by Sir Humphrey Davy in 1808
-Discovered by Carl W. Scheele in 1774
-Barium was first isolated by electrolysis of molten barium salts in 1808, by Sir Humphry Davy in England.
-Used in the Middle Ages in Alchemy
-Comes from Barite and whiterite.
-Commercially obtained from the electrolysis of molten Barium Chloride, BaCl2
-Solid at STP
-Silvery-White- oxidizes to Dark gray
-Crystal structure- Body-centered cubic
Melting point- 1000k Boiling point- 2170k Emerald Beryl Our bodies need
the correct amount
of magnesium in our diets to sleep
properly. About 13% of our planet's mass comes from magnesium. This means that there is enough magnesium in earth to make a planet with the same mass as Mars AND have enough magnesium left over to make three objects the same mass as our moon. Abundance Foods:
Spices, nuts, cereals, coffee,
cocoa, tea, fish, dairy products,
lean meat, whole grains, seeds,
and vegetables, especially green leafy
vegetables as they contain
chlorophyll. Obtained commercially
by the ‘Pidgeon’ process.
This high temperature method
uses silicon to extract magnesium
from minerals such as dolomite
or magnesite or saltwater. It is found in large deposits of
magnesite, dolomite, and
other minerals, and in
mineral waters, where
magnesium ion is soluble. Chemical Properties -Reacts with Chalcogens, oxygen, carbon, nitrogen,
phosphorus, silicon, and hydrogen, water and alcohol
-The metal is readily attacked by most acids.
-Barium combines with several metals, including Aluminum, Zinc, Lead and Tin, forming intermetallic phases and alloy.
-Oxidation State of +2 -Tube TVs- in the tubes to remove gasses
from gathering up
-Alloy with nickel- Spark Plugs
At low doses, barium ions act as a muscle stimulant Uses Hazards -Water-soluble barium compounds are poisonous.
-Higher doses affect the nervous system, causing
cardiac irregularities, tremors, weakness, anxiety, dyspnea and paralysis. 2 Interesting Facts The production of pure oxygen in the Brin process was a large-scale application of barium peroxide in the 1880s
Used in fireworks to create a Green colored blast The abundance of barium is 425bbp in the Earth's crust
-The barite reserves are estimated between 0.7 and 2 billion tonnes Abundances Isotopes 138Ba (71.7%),
132Ba (0.1%). longest Half-life
Barium 133- 10.51 years The abundance in humans is 300 ppb In the Solar System there is an abundance of 10ppb Protons: 88
Electrons:88 Line Spectra of Barium and Radium -Radium was discovered by Marie Curie and her husband Pierre on December 21, 1898.
- First isolated in 1910 by Marie Curie in the electrolysis of radium chloride.
-Radium is obtained in residues taken from uranium production.
-Radium is a white, silvery metallic , solid metal
-Solid at STP
-Highly radioactive- decays to radon gas.
- Crystal Stucture- Body- centered cubic
-Glows slight blue color
Melting point- 973k
Boiling point- 2010k Discovery Sources Physical properies Chemical Properties Behaves similarly to barium. -When exposed to air, radium reacts violently to form radium nitrate.
-Turns the metal black from white.
-Reacts with water
-Maintains its self at a higher temperature because of its high radioactivity.
-Phosphoresent Hazards -Radium is highly radioactive and hence carcinogenic
- replaces calcium in bones after longterm exposure
-Leads to cancer Uses -Radium was used in the production of luminous paints, but this is now considered too dangerous.
-Radium chloride was used medicinally to produce radon gas for cancer treatment. Safer treatments are now available. 2
interesting facts -At one point in time it was used
in toothpaste, hair creams, and in
drinking water because of its supposed
-Radium derives from the latin word for
radius or ray. Abundances -The abundance of radium in the earths crust is 0.00010 ppb
-In humans the amount is 0.000001 ppb
-In the universe the abundances is 0 ppb Isotopes 33 know Isotopes Most common-
Ra-222 38.0 seconds
Ra-223 11.43 days
Ra-224 3.66 days
Ra-225 14.9 days
Ra-226 1600.0 years
Ra-228 5.76 years Zero Stable Isotopes Discovery Sources Physical Properties