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It is abundantly clear that magnesium is in vast molar excess:
i.e.
Clearly,
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Oxygen and hydrogen react explosively to form water. In one reaction, 6 g of hydrogen combines with oxygen to form 54 g of water. How much oxygen was used?
I'll show you two approaches to solving this problem, one really short and one relatively long. $color(white)(.)$ SHORT VERSION The problem tells you that $"6 g"$ of hydrogen gas,...
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A 10.0 g sample of magnesium reacts with oxygen to form 16.6 g of magnesium oxide. How many grams of oxygen reacted?
The idea here is that chemical reactions must obey the Law of , which states that when a chemical reaction takes place, the total mass of the reactants must be...
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In the gravimetric analysis of magnesium, when combusted to magnesium oxide, why should care be taken to collect all of the magnesium oxide is retained by the crucible?
.....the following reaction occurs. $Mg(s) + 1/2O_2(g) rarr MgO(s)+Delta$ The magnesium oxide is a finely divided white solid, and CAN be lost during the exothermic combustion reaction. Nevertheless, given complete...
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Hydrogen and oxygen react chemically to form water. How much water would form if 4.8 grams of hydrogen reacted with 38.4 grams of oxygen?
We need a stoichiometric equation for water synthesis: $H_2(g) + 1/2O_2(g) rarr H_2O(l)$ Clearly, dihydrogen must be present in a 2:1 molar ratio with respect to dioxygen. $"Moles of dihydrogen"$...
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When heated in a crucible, magnesium metal reacts with oxygen from the air to form magnesium oxide in a synthesis reaction. How do you expect the mass ofthe crucible and its contents to change after the reaction?
We can represent the oxidation so.......... $Mg(s) +1/2O_2(g) stackrel(Delta)rarrMgO(s)$ And thus combustion of one equiv magnesium metal, $24.3*g$, should lead to formation of approx. $40*g$ magnesium oxide. And thus the...
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6.25 grams of pure iron are allowed to react with oxygen to form an oxide. If the product weighs 14.31 grams, find the simplest formula of the compound?
Your goal when determining the simplest formula, which is another term used for empirical formula, is to find how many moles of each element you have in your sample....
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A mixture of magnesium carbonate and magnesium oxide of mass $12.46*g$ underwent fierce heating. After heating, there were $8.43*g$ of salt remaining. What were the masses of magnesium carbonate, and magnesium oxide in the original sample?
$MgCO_3(s)+DeltararrMgO(s) + CO_2(g)$ And we would have to heat the magnesium salt VERY fiercely to get complete decarboxylation. $"Moles of carbon dioxide LOST"$ $=$ $(4.40*g)/(44.01*g*mol^-1)$ $=0.010*mol$ And therefore in...
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Solid bismuth oxide can react with carbon to form bismuth metal and carbon monoxide. How many grams of bismuth oxide reacted if 60.7 grams of bismuth is formed?
The answer is $67.6$ $g$. First, start with the balanced chemical equation $Bi_2O_(3(s)) + 3C_((s)) -> 2Bi_((s)) + 3CO_((g))$ We know the molar masses of $Bi$ ($209g/(mol)$) and $Bi_2O_3$ ($466g/(mol)$),...
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How many moles of magnesium oxide are formed when 4 moles of magnesium react with oxygen in the reaction $2Mg + O_2 -> 2MgO$?
Let us first look at the equation. As per the equation , 2 moles of Magnesium produces two moles of MgO . If we double the amount of Magnesium ,...
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Magnesium will burn in oxygen to form magnesium oxide as represented by the following equation
2Mg(s) + O2(g) -> 2MgO(s)
What mass of magnesium will react with 498.1 ml of oxygen at 145.2 degrees Celsius and 0.927 atm?
We use the Ideal Gas Equation to access the molar quantity of dioxygen gas............ $n=(PV)/(RT)=(0.927*atmxx0.4981*L)/(0.0821*(L*atm)/(K*mol)xx418.4*K)=1.34xx10^-2*mol$. Given the stoichiometric equation.......... $Mg(s) + O_2(g) rarr MgO(s)$ There are $1.34xx10^-2*molxx24.31*g*mol^-1=??*g$ metal required for...
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